Things to Know Chem 30A 7 th Ed McMurry Chapters 1-11

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1 Things to Know Chem 30A 7 th Ed McMurry Chapters 1-11 Chapter 1 1. What is matter? 2. Be able to identify physical vs. chemical properties (what s the difference?) and physical vs. chemical changes 3. Be able to explain the differences between solids, liquids and gases both their observable properties and their differences at the molecular level. 4. Be able to define and explain the differences between the categories of matter: pure substances, mixtures, elements, compounds, homogeneous mixtures, heterogeneous mixtures. How can mixtures be separated? How can compounds be separated? Can elements be separated? 5. Memorize the names and symbols of the elements on the handout. Given the symbol, write the name of the element (with correct spelling) and vice versa. 6. Chemical reactions: reactants products 7. What are the metric units of mass, length, and volume? 8. Know these metric prefixes and how to use them as conversion factors: kilo, centi, milli, micro, nano. 1 kg = 1000 g 1 g = 100 cg 1 g = 1000 mg 1 g = 10 6 µg 1 g = 10 9 ng Difference between mass and weight 9. 1 cm 3 = 1 ml (memorize) 10. Any measurement involves some uncertainty. Be able to state the amount of uncertainty in a measurement and the number of sig figs (Example: 15.6 ml has an uncertainty of ± 0.1 ml and 3 significant figures.) Be able to state the range of possible values. (A measurement of 15.6 ml could range from 15.5 ml to 15.7 ml.) 11. Know how to interpolate a measurement (estimate between the lines) 12. Be able to round results of calculations to the correct number of significant figures. Remember, the rule for multiplying and dividing is different from the rule for adding and subtracting. 13. Multiplying and dividing: lowest number of sig figs. 14. Adding and subtracting: lowest number of decimal places and/or greatest uncertainty. 15. Be able to write numbers in scientific notation or given a number in scientific notation, write it as a regular number. 16. ***Unit conversion problems! (Do lots of practice!) In a problem, when they state a relationship between two things, you can write a conversion factor representing that relationship. For example: each pill contains 500 mg of the 1

2 active ingredient : the conversion factor would be (1 pill/500 mg) OR (500 mg/1 pill). Expect many unit conversion problems. 17. Always ask yourself whether your answer makes sense. 18. Be able to convert C to K and K to C (Remember to add or subtract 273.) Kelvin temperature is always larger! 19. Be able to convert C to F or vice versa. (You will be given one equation, and you might need to rearrange it algebraically.) 20. Be able to use the equation q = sm T. Know what each of the letters stands for. Be able to solve for ANY of these four variables. The problem might also involve some unit conversions. 21. Be able to do the calculation from lab 7. (Finding the specific heat of a substance using a calorimeter.) 22. Definition of density, how to calculate density, **how to use density as a conversion factor!!** (You will be using density as a conversion factor frequently.) 23. What is specific gravity? Chapter 2 1. Postulates of Atomic Theory. What does a chemical reaction involve (in terms of the atoms)? 2. Subatomic particles: proton, electron, neutron. Know charge, mass (atomic scale), location of each 3. Structure of atom 4. Atomic number, mass number what are they? 5. What are isotopes? 6. Be able to write or interpret an isotope symbol. (How many protons, electrons, and neutrons?) 7. How is the atomic mass listed on the periodic table related to the mass numbers of the elements? 8. Know how to calculate the average atomic mass of an element, given the masses and percent abundances of each isotope. 9. How is the periodic table arranged? 10. Terms: period, group, main-group elements, transition metals, inner transition metals. Know what these mean and give examples of each. 11. Properties of metals, nonmetals, metalloids 12. Alkali metals, alkaline earth metals, halogens, noble gases: What are they? What are their general properties? 13. What does it mean if something is quantized? Give an example of something that is quantized and something that isn t quantized. 14. For shells 1, 2, 3, and 4: how many electrons can they hold? How many subshells do they contain? 15. What is an orbital? 16. Higher shell number means electrons have higher energy, the orbitals are larger in size, electrons are further from the nucleus on average, etc. (compared to a lower shell number) 17. Shapes of s, p, and d orbitals (sketch) 2

3 18. Do we know anything about the path an electron follows? Explain. 19. Be able to write the electron configuration for any element (up to Ba). 20. Be able to write the noble gas notation (abbreviated electron configuration) for any element up to Ba. 21. Be able to draw the arrow diagram for any element (up to Ba). Be able to draw the arrow diagram for the outer electrons of any element (up to Ba). 22. What are valence electrons? How can you tell the number of valence electrons of a main-group element? 23. Write electron-dot symbols for atoms (valence electrons are shown as dots. Don t pair them until necessary). Chapter 3 1. What is an ion? A cation? An anion? 2. Metals tend to form cations, nonmetals tend to form anions. 3. What is the significance of a filled shell of electrons? 4. Be able to write the electron configuration for a main-group monatomic ion. (Determine the number of electrons, and fill orbitals in order. A main-group monatomic ion will have a filled shell electron configuration.) 5. Predict the charge for any main-group monatomic ion. 6. What is an ionic compound? What holds ionic compounds together? 7. What are some typical properties of ionic compounds? 8. Transition metal ions can have more than one possible charge. 9. Names of ions: Na + sodium ion, Fe 3+ iron (III) ion, Cl - chloride ion. (cation name is the same as the corresponding element, transition metals must include a roman numeral that indicates charge, and monatomic anions end in ide.) 10. Polyatomic ions: memorize names, formulas, and charges. (handout) 11. To write the formula of an ionic compound, you need to know the charges on the ions it contains. How many of each ion are needed to make an uncharged compound? (Na + and Cl - make NaCl; K + and O 2- make K 2O; Fe 3+ and O 2- make Fe 2O 3; Mg 2+ and O 2- make MgO; Ba 2+ and PO 4 3- make Ba 3(PO 4) 2, etc.) 12. Naming ionic compounds: name cation, name anion. No need to explicitly say how many of each ion there are. (Examples: NaCl is sodium chloride, K 2O is potassium oxide, Fe 2O 3 is iron (III) oxide, MgO is magnesium oxide, Ba 3(PO 4) 2 is barium phosphate.) 13. Be able to name any ionic compound. If you are given the name, be able to write the correct formula for any ionic compound. 14. Know the general properties of acids and bases. Be able to recognize acids and bases. Chapter 4 1. Nonmetals tend to share electrons. 2. What is a covalent bond? 3. Which elements are diatomic? 3

4 4. What is the usual number of bonds made by atoms in groups 4, 5, 6, and 7? How many lone pairs will each tend to form? 5. Are there exceptions to the above tendencies? 6. What is the difference between a single, a double, and a triple bond? 7. Be able to draw the Lewis structure of any molecular compound or polyatomic ion. 8. If the molecule contains carbon and H, N, O, or halogens, you can just figure out the Lewis structure based on the normal number of bonds each element tends to make. 9. In many cases, you will need to use the systematic method for determining Lewis structures. First, determine the skeleton structure: how are the atoms connected? Then, count up the total number of valence electrons in your molecule. Your final structure must have this number of valence electrons: no more, no less! Then, fill in octets for the outer atoms first. Put any remaining electrons on the central atom. If is seems like there aren t enough electrons, form multiple bonds. Make sure to satisfy the octet rule for all atoms! 10. After drawing the Lewis structure, be able to determine the shape and the bond angles for a molecule or polyatomic ion. If the molecule has more than one central atom, state the shape/geometry around each central atom. 11. What are the differences between ionic, polar covalent, and nonpolar covalent bonds? (sharing of electrons, difference in electronegativity) 12. Know the general trend in electronegativity values according to the periodic table. 13. Predict whether given bonds are polar or nonpolar (with electronegativity values or in some cases without them). 14. Be able to determine whether a molecule or polyatomic ion is polar (not symmetric) or nonpolar (symmetric). 15. Naming binary molecular compounds: use prefixes. CO 2 is carbon dioxide, CO is carbon monoxide, N 2O is dinitrogen monoxide. Be able to name or write the formula for any binary molecular compound. 16. What are some differences between molecular and ionic compounds? Chapter 5 1. Be able to balance chemical reactions. 2. Be able to predict the products of a precipitation reaction. (You must first determine the formulas and charges of the ions present. Switch the ions and write the formulas of the new compounds. Check solubility rules to get phases, and balance the equation.) 3. Be able to write the balanced equation for a precipitation reaction. 4. Be able to write the balanced equation for an acid-base reaction. 5. Be able to write the balanced equation for the reaction between a carbonate or bicarbonate compound and an acid. The products are water, carbon dioxide gas, and an ionic compound. 6. Given a simple oxidation-reduction reaction, determine what substance is oxidized, what substance is reduced, the oxidizing agent, and the reducing agent. Be able to explain your reasoning. (Determine the charge of each atom 4

5 before and after the reaction. Look at whether the charge is increasing or decreasing. Single elements by themselves with no charge written have a zero charge. For ionic compounds, determine the charges on each ion.) 7. Be able to write total ionic and net ionic equations for any precipitation or acidbase reaction. (Any soluble ionic substances or strong acids must be written as separate ions.) Chapter 6 1. What is a mole? 2. Be able to calculate molar mass of any given molecule. 3. Be able to convert grams moles or moles grams. 4. Be able to convert # moles # molecules or # molecules #moles. 5. Be able to convert # grams molecules or the reverse. 6. Be able to convert # moles of one substance to # moles of another substance. You need a balanced equation. 7. Be able to convert grams of one substance to grams of another substance. (g of X mol X mol Y g of Y) You will need a balanced equation and the molar masses of X and Y. 8. Be able to convert g of X moles of Y or moles of X g of Y 9. Be able to do variations of these stoichiometry problems. These variations could involve density or number of molecules or other conversions. 10. Given amounts of two reactants, determine which is the limiting reactant and determine the amount of product that will be formed. Remember that the limiting reactant gets used up first. When it gets used up, the reaction stops. Approach to these problems: do two stoichiometry calculations: Convert mass of the first reactant to mass of product. Convert the mass of the second reactant to the mass of product. Which reactant produces the smallest amount of product? That s the limiting reactant. The amount of product formed will be the smaller amount! (Do not add the masses from each.) 11. Be able to calculate percent yield for a reaction. 12. What is the difference between theoretical yield, actual yield, and percent yield? Chapter 7 1. State the law of conservation of energy. 2. What is the H of a reaction? 3. Exothermic reaction: what is the sign of H? What is more stable, reactants or products? What contains stronger bonds, reactants or products? Draw an energy diagram. 4. Answer the above questions for an endothermic reaction. 5. The amount of heat a reaction gives off or absorbs depends on how much of the reactants are present. More stuff present means more heat involved. 6. A + 2 B 3 C + D H = kcal means: 350 kcal given off per 1 mole of A, per 2 moles of B, per 3 moles of C, or per 1 mole of D. 5

6 7. Be able to calculate the energy given off or absorbed by a specific amount of a chemical, given a reaction and its H value. 8. If you reverse a reaction, the sign of H changes. 9. Define spontaneous and nonspontaneous. 10. What two factors affect whether or not a reaction is spontaneous? 11. What is entropy? What does it mean if S is positive? What if S is negative? 12. Be able to predict the sign of S for a reaction or process. 13. What is G? Define endergonic and exergonic. 14. G = H T S 15. What two things must happen for molecules to react? 16. Define/explain the activation energy of a reaction. 17. Be able to draw an energy diagram for a reaction, indicating the activation energy on the graph. 18. How does the size of activation energy affect the rate of the reaction? 19. What factors affect reaction rates? Briefly explain each. 20. What is meant if a reaction is at equilibrium? Does it mean the amounts of reactants and products are equal? 21. Be able to write the equilibrium constant expression for any reaction. 22. Understand the meaning of the size of K. If K is very large, what does that mean about the relative amounts of reactants and products at equilibrium? What if K is small? What if K is medium-sized? Be able to explain. 23. What is Le Chatelier s principle? 24. Be able to explain what will happen to a system at equilibrium if you add or remove one of the reactants or products. 25. Be able to explain what will happen to a system at equilibrium if you increase or decrease the temperature. (You will need to know the sign of H.) 26. Be able to explain what will happen to a system at equilibrium if you increase or decrease the pressure. (Look at number of moles of gas on each side of the equation.) Chapter 8 1. Know names of phase changes and what they correspond to: melting, boiling, evaporating, condensing, and freezing. Which ones require an input of heat and which ones release heat? Know what bp and mp correspond to. 2. List the assumptions of kinetic-molecular theory. 3. Why do gases exert pressure on the walls of their container? (Explain at the molecular level) 4. What is pressure? What are units of pressure? How is atmospheric pressure measured? How does atmospheric pressure vary? 5. Boyle s Law: P 1V 1 = P 2V 2 As pressure increases, volume decreases. (Used if n and T are constant.) Why? Explain at the molecular level. 6. Charles Law: as T increases, volume also increases. V is proportional to temp in V 1 T 1 = V 2 T 2 6

7 Kelvin. (Used if n and P are constant.) Why? Explain at the molecular level. 7. Gay-Lussac s Law: as T increases, P increases (assuming n and V are constant). P 1 = P 2 T 1 T 2 Why? Explain at the molecular level. 8. Combined gas law (used when n is constant but P, V and T are changing) P 1 V 1 = P V 2 2 T 1 T 2 9. Avogadro s law: volume is proportional to moles of gas (assume P and T are constant). Why? Explain at the molecular level. V 1 = V 2 n 1 n What is STP? What is the molar volume of an ideal gas at STP? 11. Be able to use the molar volume of an ideal gas as a conversion factor. (Molar volume of a gas at STP is 22.4 L/mole.) 12. Ideal gas law: PV = nrt. R = L atm/k mol. Be able to solve for any of the variables if you know the other ones. Make sure that units of everything match the units of R. 13. For mixtures of gases, each exerts its own pressure. P total = P 1 + P 2 + P Partial pressure of X = (fraction of X)(total pressure) 15. Intermolecular forces: in general, the stronger the IMF s, the higher the boiling point, the lower the vapor pressure, and the higher the heat of vaporization. Why? Strong IMF s mean that the molecules tend to stick together and that it s hard to separate them. 16. What are dipole-dipole forces? What type of molecules has dipole-dipole forces? 17. What are London dispersion forces? What factors affect the strength of London forces? What types of molecules have London forces? 18. What is hydrogen bonding? What type of molecules can form hydrogen bonds? Remember that the H and the two atoms on each side of the H must be in a straight line. 19. Given the formula of a molecule, be able to identify the types of IMF s present. 20. Given several formulas of molecules, be able to reason out which one would have the strongest intermolecular forces and therefore the highest boiling point, highest heat of vaporization, and lowest vapor pressure. Be able to rank the compounds in order of boiling point and explain your reasoning thoroughly. 21. What is vapor pressure? What does it depend on? At what point does a liquid begin to boil? 22. What is the normal boiling point of a liquid? Can the boiling point ever be different from the normal bp? Explain. 23. Explain how boiling point varies with altitude. 24. Explain how a pressure cooker works. 25. List some unusual properties of water. 26. What is the difference between a crystalline solid and an amorphous solid? 27. H fus, H vap: know how to use them to calculate the amount of heat needed to melt or vaporize a substance. 28. Phase changes occur at a constant temperature. 7

8 29. Be able to calculate the energy needed for a combination of phase changes and temperature changes. For a temperature change, use q = sm T. Chapter 9 1. Define: solution, solvent, solute. 2. Can solutions be gases? Can they be solids? 3. Solubility Like dissolves like look for similar types of IMF s. Polar + polar mix, nonpolar + nonpolar mix, but polar + nonpolar don t mix. 4. What types of compounds generally dissolve in water? 5. What types of compounds generally don t dissolve in water? 6. Some molecules have a region that is nonpolar and a region that is polar. If most of the molecule is hydrocarbon, it s probably not soluble in water overall. (Example: C 8H 17OH has a long nonpolar hydrocarbon section, but it can also hydrogen bond. However, the nonpolar part is very long, so it is mostly nonpolar. This compound is not soluble in water, even though it can hydrogen bond.) 7. For organic molecules containing OH groups, look at the ratio of carbons to OH groups in the molecule. Higher C/OH ratio is less soluble in water. Be able to explain why. 8. Given some compounds, be able to explain which ones would be more soluble in water and which ones would be more soluble in a nonpolar solvent, and be able to explain why for each. 9. Definition of miscible 10. The solubility of a substance is usually expressed in units of # g of substance that dissolves in 100 ml of water at a specified temperature. The solubility expressed in this way is the maximum that can dissolve under those conditions. 11. Differences between saturated, unsaturated, and supersaturated solutions. Which one is unstable? 12. If you are given the solubility of a compound in g solute/100 ml water, and you are given the number of grams of solute combined with a specified amount of water, be able to determine if the solution is saturated or unsaturated. 13. The solubility of a substance depends on temperature. Most solids get more soluble as temperature increases. Gases, however, get less soluble as temp increases. Describe an application of this. (See section 9.5) 14. How does the solubility of a gas depend on the pressure of a gas over the liquid? Describe an application. (Sec. 9.6) 15. Units of concentration: for each, be able to calculate the concentration in the appropriate units. Units to know: M, % (m/m), % (m/v), % (v/v), ppm, ppb. 16. Be able to rewrite each of these concentration units as a conversion factor and use it in calculations. Units to know: M, % (m/m), % (m/v), % (v/v), ppm, ppb. 17. Be able to do stoichiometry problems involving molarity. (Use molarity as a conversion factor.) Remember, V M = moles of solute, as long as volume is in units of liters. 18. Dilutions: use M 1V 1 = M 2V 2. In this equation, V 2 is the total final volume. V 2 = V 1 + V added. You can also use c 1V 1 = c 2V 2 if the concentration is not in molarity. 8

9 19. Explain the differences between strong, weak, and nonelectrolytes. Give an example of each. 20. Define an equivalent of an electrolyte. Be able to calculate the mass of one equivalent of an electrolyte. Be able to use the number of meq/l as a conversion factor in a calculation. Also, 1 Eq = 1000 meq. 21. Osmosis what is it? 22. Decide which direction water will flow across a semipermeable membrane. Explain the reason for the direction of water flow. 23. What is osmolarity? What is the signficance of osmolarity? 24. Decide whether cells will shrink or swell in various solutions. Chapter Arrhenius definition of acids and bases 2. Know the formulas and names of the acids and bases in section Brønsted-Lowry definition of acids and bases: proton donor/proton acceptor. 4. Given an equation for a proton-transfer reaction, label the acids and the bases. 5. Water can act as an acid or a base in proton-transfer reactions. 6. What is a conjugate acid-base pair? 7. Given a formula, write the formula of its conjugate acid or base. 8. Be able to write the molecular equations for the following types of reactions: Acid + base water + a salt Acid + a carbonate or bicarbonate compound water + CO 2 + a salt Acid + ammonia an ammonium salt 9. There are only seven strong acids, and the rest are weak acids. Know these strong acids: HCl, HBr, HI, HNO 3, H 2SO Are all weak acids equally weak? 11. The higher the K a, the stronger the acid. Why? 12. Write the equation for the dissociation of a weak acid in water. 13. Write the equilibrium-constant expression for the dissociation of a weak acid in water. 14. Write the equation for the autoionization of water. 15. In any aqueous solution (at 25 C), [H 3O + ][OH - ] = K w = Given hydronium ion concentration, calculate hydroxide concentration (or vice versa) 17. How do you know if a solution is acidic or basic? 18. As hydronium ion increases, what happens to the corresponding ph? 19. Given hydronium ion concentration, calculate ph on your calculator. 20. Estimate approximate ph from the exponent of the hydronium concentration. 21. Given the ph, calculate hydronium concentration (using your calculator) and be able to estimate it. 22. Based on the ph, how do you know if a solution is acidic or basic? 23. Describe what a buffer, what it is made of, and what it does. 24. How does a buffer resist ph changes? 25. Write a reaction for what happens when an acid or a base is added to a specific buffer. (Any added acid will react with the conjugate base in the buffer. Any added base will react with the conjugate acid in the buffer.) 9

10 26. Titration problems: Need to write a balanced equation for the reaction. Do not use the dilution formula!! Start with the substance that you know the most about, and find moles of that substance. Convert to moles of the other substance. Finish the problem, either calculating the volume of solution needed or the molarity of the solution. Chapter Know how to write a nuclide symbol and know what it means. 2. What is a nucleon? 3. What are some unusual characteristics of nuclear reactions? 4. Know the symbols for alpha, beta, and gamma radiation. 5. Know the penetrating properties of each of the above types of radiation. 6. What does it mean if an isotope is radioactive? 7. Which elements are artificial? Which elements are radioactive? Are there more stable or more radioactive isotopes? 8. What is nuclear decay? 9. Be able to write or complete nuclear equations. Remember that total mass number and total atomic number must be conserved. 10. Be able to write the nuclear equation for alpha, beta, or positron emission, given a nuclide symbol. 11. What is meant by the term half-life of a radioisotope? 12. Determine how much of a sample is left after a certain amount of time has passed. Determine how much time it will take for the amount of radioactivity to fall to a certain level. (Make a chart) 13. What is ionizing radiation? What does it do? 14. What are some devices used to detect radiation? 15. Know the definitions of these units of radiation: curie (Ci), rad, rem 16. What is background radiation? What are some sources of background radiation? 17. Define LD What is an artificial transmutation? 19. Explain the differences between fission and fusion. Why do both of them give off a lot of energy, even though they seem to be opposites? 20. What is the general idea behind radiocarbon dating? How can you determine the age of an ancient object that was once alive? 10