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SGSF-1K PreAP ESPS Spring Final Study Guide KEY Unit 1 – Safety & Matter 1) Why should you clean any dirty glassware before using it for an experiment? Avoid contamination 2) Why do goggles or face shields provide more eye protection than personal eyeglasses? Eyeglasses don’t protect from side splashes. 3) What should you do with long hair in the laboratory? Tie it up. 4) What can happen if you wear loose clothing or jewelry in a laboratory? Catch on things, get contaminated by chemicals 5) What is the best type of footwear to wear in a lab? Closed toed shoes 6) What should you do if a piece of lab equipment is not working properly? Tell the instructor 7) What should you do if acid is splashed on your skin? What if it’s splashed all over your clothes? Skin: Flush with lots of water; Clothes: get them off and flush skin with lots of water 8) Why should you clean your hands after working with chemicals? What should you use to clean them? Chemicals could be on your hands – contamination. Wash with soap and water 9) How does hot glass look compared to cold glass? Same 10) What does it mean if a material is “flammable”? It will easily ignite and burn. 11) What should you do if an uncontrolled fire occurs in the laboratory? Immediately tell the instructor 12) When should you wear special eye protection devices (such as goggles) in the laboratory? Whenever working with chemicals or glassware 13) In the lab, what should you do it you don’t understand a direction or part of a procedure? Ask the instructor 14) What are you supposed to do with chemical wastes after you finish a lab? Dispose of according to instructions from instructor 15) What could happen if you point the end of a test tube that you’re heating towards someone? Material could shoot out the end and hit someone. 16) What’s the best way to pick up a piece of glassware you’ve been heating? With tongs 17) If you get injured in the laboratory (cut, burn, etc.), what should you do? Immediately tell the instructor 18) Why is it important to check glassware for chips or cracks? Cracks and chips weaken glass so it more easily breaks 19) Is it appropriate to return all unused chemicals to their original containers? No 20) Should you start working on a lab even if the instructor is not yet present? No 21) When is okay to remove chemicals or other equipment from the laboratory? Never (only under direct supervision of an instructor) 22) Why are unauthorized experiments prohibited? Potential for unexpected reactions and/or consequences 23) When are students allowed to enter the chemical preparation/storage area? Never 24) Why do people wear laboratory aprons? To protect clothes and skin from chemical splashes/contamination 25) Is it okay to pick up broken glass with your bare hands as long as the glass is placed in the trash? No, ask instructor (who will use tongs and broom with dust pan) 26) Why should you never leave a lit burner unattended? Flame could go out and allow unburned gas to accumulate in room – potential for explosions 27) What instrument is used to weigh objects? Balance 28) What piece of lab equipment is a small glass container used to view chemical reactions or to heat small amounts of a substance? Test Tube 29) What instrument is used to measure volume very precisely? Graduated Cylinder 30) What instrument is a wide-mouthed container used to transport, heat, and store substances? Beaker 31) Define: Matter – ANYTHING THAT HAS MASS AND TAKES UP SPACE Substance – MATTER THAT ALWAYS HAS THE EXACT SAME COMPOSITION Mixture – 2 OR MORE SUBSTANCES TOGETHER - COMPOSITION CHANGES Element – SUBSTANCE THAT CAN’T BE BROKEN DOWN INTO SIMPLER SUBSTANCES Compound – COMPOSED OF 2 OR MORE ELEMENTS IN FIXED RATIO 32) Name two examples each of an element, and a compound Element: OXYGEN, CARBON, GOLD, SILVER, ETC. Compound: WATER, SALT, SUGAR, CARBON DIOXIDE, SILICON DIOXIDE, ETC. 33) How are a compound and a mixture different? MIXTURE RETAINS MANY PROPERTIES OF IT THE COMPONENTS; NO FIXED RATIO 34) Describe and give an example of each: Heterogeneous mixture: SAND, RIVER WATER, SALSA, CHOCOLATE CHIP COOKIE Homogeneous mixtures. MILK, KOOL-AID, SUGAR COOKIE, GASOLINE Record which picture represents each type: a - HOMOGENEROUS; b - HETEROGENEOUS 35) Define: Solution – HOMOGENOUS MIXTURE OF SOLUTE DISSOLVED IN SOLVENT o Solvent – DOES THE DISSOLVING o Solute – IS DISSOLVED Suspension – HETEROGENEOUS MIXTURE THAT SEPARATES INTO LAYERS OVER TIME Colloid – HOMOGENEOUS MIXTURE WITH PARTICLE SIZE BETWEEN SUSPENSION AND SOLUTION Alloy – HOMOGENEOUS MIXTURE OF 2 OR MORE METALS 36) Describe the following: (in terms of motion, volume & shape) Solid – Vibrate - Locked in place; FIXED VOLUME; FIXED SHAPE Liquid – Glide – still attracted to each other; FIXED VOLUME; INDEFINITE SHAPE Gas – Bounce; INDEFINITE VOLUME; INDEFINITE SHAPE 37) Describe these changes as physical or chemical. Substances are…… Mixed P Burned C Melted P Rusted C Cut P Cooked C Spoiled C Dissolved P Digested C Bent P Frozen P Reacted C 38) What does the Kinetic Theory state? ALL PARTICLES OF MATTER ARE IN CONSTANT MOTION 39) The temperature of a substance does not change during a PHASE CHANGE 40) Complete this diagram with the 6 common phase changes: A. DEPOSITION B. SUBLIMATION C. VAPORIZATION D. CONDENSATION E. MELTING F. FREEZING 41) Define a. Heat of Fusion – ENERGY THAT MUST BE ABSORBED TO CHANGE FROM SOLID TO LIQUID b. Heat of Vaporization – ENERGY THAT MUST BE ABSORBED TO CHANGE FROM LIQUID TO GAS 42) Label the following phase change diagram with the states of matter and all 6 phase changes. Use arrows to display the direction of each specific phase change. 43) What volume of silver metal will have a mass of exactly 2500.0 g. The density of silver is 10.5 g/cm3. D = M/V V = M/D V = 2500.0g/10.5g/cm3 = 238.1cm3 238cm3 44) A block of lead has dimensions of 4.50 cm by 5.20 cm by 6.00 cm. The block weighs 1587 g. From this information, calculate the density of lead. D = 1587g/(4.50cm x 5.2cm x 6.00cm) = 11.3g/cm3 45) What is the mass of the ethanol that exactly fills a 200.0 mL container? The density of ethanol is 0.789 g/mL. M = D x V M = (0.789 g/mL)(200.0mL) = 157.8g 158g Unit 2 – Atomic Structure & Periodic Table 46) Subatomic particle descriptions: A. Name the three subatomic particles: proton, neutron, electron B. Give the location where each can be found: nucleus – proton & neutron; electrons in electron cloud C. Give their electric charges: proton is +; electron is - ; neutron is neutral D. Give their relative masses: proton ~ neutron are about equal in mass (1); electron ~1/2000th as much E. Describe the composition and characteristics of the nucleus: protons with (+) and neutrons (neutral) for net + charge; almost all of atom’s mass in nucleus 47) Atomic number and mass number: A. What determines the atomic number of an atom? - # of protons B. What determines the atomic mass of an atom? - # of protons + # of neutrons C. How can you determine how many neutrons will be in a given atom? Atomic mass # – atomic number (i.e. # of protons) = # of neutrons D. For an atom to be neutral, what subatomic particles have to have been present in the same number? – protons (+) and electrons (-) E. What number is unique for any given element? Protons (atomic #) 48) Isotopes & Ions A. Define “isotope”: - atoms of an element with differing numbers of neutrons (and hence, differing atomic mass) B. What is same about all isotopes of a given element? # of protons (& # of electrons in neutral atom) C. What is different between isotopes of a given element? # of neutrons D. What determines the listed atomic mass for an element with many isotopes? - weighted average, isotopes have a greater effect than uncommon ones E. Define “ion” – an atom of an element that carries a charge (not neutral) F. How does a neutral atom become an ion? – gaining of losing electrons (gaining = negative charge; losing = positive charge) G. How do you determine the charge of an atom? - # of protons - #of electrons 49) Electron energy levels: A. The number of energy levels filled in an atom is determined by what? What is the new # sequence? – how many electrons the atom has (always fill lowest first); 2, 8, 18, 32 B. What causes an electron to jump to a different energy level? Atom gains or loses energy C. Define ground state: - all the electrons in an atom have lowest possible energy D. Define excited state; - at least one electron at a higher energy level than ground state E. What causes the glow of a neon light? – electrons dropping from excited state to ground state by emitting energy as light F. In a flame test, which color corresponds with the highest excited energy level and which with the lowest? Red is lowest; violet is highest 50) Complete the following table: Element Symbol Se C P Fe Xe Zr Br Ca Mn Na Ag Atomic Number 34 6 15 26 54 40 35 20 25 11 47 Atomic Mass 78.96 12.011 30.974 55.847 131.30 91.22 79.904 40.08 54.938 22.990 107.87 Mass Number 79 12 31 56 131 91 80 40 55 23 108 Protons Neutrons Electrons 34 6 15 26 54 40 35 20 25 11 47 45 6 16 30 77 51 45 20 30 12 61 34 6 15 26 54 40 35 20 25 11 47 51) Draw a Bohr Diagram for Carbon, Sodium, and Sulfur. C 52) Calculate the average atomic mass for the following: A. Calculate the average atomic mass of sulfur if 95.00% of all sulfur atoms have a mass of 31.972 amu, 0.76% has a mass of 32.971amu and 4.22% have a mass of 33.967amu. 32.021amu B. An atom has the following three isotopes: 24 amu (percent abundance = 78.99%), 25 amu (percent abundance = 10.00%), and 26 amu (percent abundance = 11.01%). Calculate the average atomic mass. 24.320amu 53) History of the Periodic Table A. Who first arranged the periodic table by atomic mass? Dmitri Mendeleev B. What characteristics did Mendeleev use to place the elements in order when creating the periodic table? Rows - Increasing mass Columns – Similar properties C. What are valence electrons? Electrons in the highest occupied energy level D. What do elements that belong to the same group have in common? They have the same number of valence electrons; similar chemical properties E. What is the “octet rule”? Atoms are most stable if they have filled or empty outer shell of electrons Filled shell contains 8 electrons (octet) Except for H and He (atomic #1 & #2) Atoms gain, lose, or share electrons to make filled or empty outer shell Atoms gain, lose, or share electrons based on what is easiest. 54) Structure of the Periodic Table A. What is the definition of atomic mass? Individual Atoms: Number of protons plus number of neutrons. (p+ + n) Elements: Weighted average of the masses of the isotopes. B. What units do you put on atomic mass? amu (Atomic Mass Units) 1 amu = 1/12 mass of Carbon-12 C. Does the atomic number increase or decrease as you move from left to right? How much does it change from one element to the next? Increases by exactly 1 for each element as you move from left to right D. The vertical columns on the Periodic Table are called what? Groups or Families (memory aid: families stand up for each other) E. The horizontal rows are called what? Periods (memory aid: read the row like a sentence and end it with a “period”.) F. What states of matter are represented by the metals at room temperature (solid, liquid, and/or gas)? Solid and liquid (all but mercury are solids) G. Every element in the Carbon group (family) has how many valence electrons? 4 H. The majority of the Periodic Table is made up of what class of element? metals I. Which group is the least likely to react? Noble Gases J. The chemical and physical properties of Li are most similar to the chemical and physical properties of what other elements? K, Rb, Cs, Fr 55) Trends on the Periodic Table A. What is the trend in reactivity within the alkali metals and the alkaline earth metals? Reactivity increases as you go down the column (most reactive elements at the bottom) B. What is the trend in reactivity within the halogens? Reactivity decreases as you go down the column (most reactive elements at the top) C. There are two groups (families) on the Periodic Table that are considered the MOST reactive, what are they (give name and number)? Group 1 (IA) – Alkali Metals Groups 17 (VIIA) – Halogens D. Using your answer from the previous question, how many valence electrons does each of those groups (families) have? Group 1 (IA) – Alkali Metals has 1 valence electron Groups 17 (VIIA) – Halogens has 7 valence electrons E. Na is a very reactive element in Period 3, what element in Period 2 would be similar in its reactivity? Lithium (both are in group 1 – IA) F. Looking at Group 2 (IIA) which element would be the MOST reactive? Radium G. If a reaction is going to take place, which Halogen would you expect to react the fastest? Fluorine H. If K reacts very violently in water, what you expect Fr to do? React even more violently I. Define Ionization Energy. Which metal has the highest ionization energy? Why? The amount of energy needed to remove an electron from an atom. Li – because it’s one valence electron is in energy level 2 which is close to the nucleus resulting in a much stronger magnetic pull on it than on the valence electrons of other members of the group which as in higher energy level. J. Define Electronegativity? Which element has the highest electronegativity? Why? The ability of an atom to attract or gain an electron F – because it only needs to gain 1 electron to fulfill the octet rule and its highest energy level is 2 so the magnetic pull on that electron will be greater due to the energy level’s proximity to the nucleus K. List the following elements in order of increasing ionization energy: Li, Ti, Ba, K, Fr Fr, Ti, Ba, K, Li L. List the following elements in order of increasing reactivity: C, S, I, Te, F C, Te, S, I F M. Label the numbers and letters in the above Periodic Table 1 – Alkali Metals 2 – Transition Metals 3 - Halogens 4 – Lanthanide Series 5 – Actinide Series 6 - Alkaline Earth Metals 7 – Noble gasses A – Metals (includes H) B – Non-metals C – Metalloids N. Use the periodic table to complete the following table: Element # of valence electrons Magnesium 2 Oxygen 6 Neon 8 Aluminum 3 Lewis Dot Diagram O. List and illustrate the 4 periodic trends we discussed in this section: Metal Reactivity Increases: Nonmetal Reactivity Increases: 56) Electron Configuration: Unit 3 – Bonding Review Ionic Bonds 57) What two types of elements will transfer electrons to form an ionic bond? Metal & nonmetal 58) What is an ionic bond? An electrostatic force that holds together a cation and anion 59) How can you describe the electrical charge of an ionic compound? Neutral 60) Metals will lose all of their valence electrons to form a positively charged cation. The oxidation number will be equal to the number of valence electrons. 61) Nonmetals will gain valence electrons to form a negatively charged anion. The oxidation number will be equal to 8 – the # of valence electrons. 62) Which elements are least likely to undergo bonding? Noble Gases 63) Write the oxidation number for the following elements. a. Cesium ___+1____ b. Potassium ___+1____ c. Aluminum ___+3____ d. Iron (III) ___+3____ e. Magnesium ___+2____ f. Barium ___+2____ g. Nitrogen ___-3____ h. Iodine ___-1____ i. Oxygen ___-2____ j. Selenium ___-2____ 64) Which alkali metal has the highest ionization energy? Hydrogen 65) Which alkaline earth metal has the lowest ionization energy? Radium 66) What properties characterize ionic compounds? High melting point & boiling point Solid at room temperature Brittle Ionic solids can’t conduct electricity Molten ionic compounds can conduct electricity Transferring of Electrons 67) Please draw the transfer of electrons for the following: a. NaBr b. MgO c. CaF2 d. K2S 68) How do the following atoms fulfill the octet rule when they form ionic bonds? What valence or oxidation number does each of these have? a. Lithium Loses 1 e-; valence 1+ b. Beryllium Loses 2 e-; valence 2+ c. Boron Loses 3 e-; valence 3+ d. Carbon Loses 4 e-; valence 4e. Nitrogen Gains 3 e-; valence 3f. Oxygen Gains 2 e-; valence 2g. Fluorine Gains 1 e-; valence 1h. Neon Full shell – no change 69) Define the indicators (what you look for) that identify each type of compound based on their names and their formulas: a. Simple Binary Ionic – i. Name – Group 1A/2A metal or Aluminum & Nonmetal with “ide” ending ii. Formula – Group 1A/2A metal or Aluminum & Nonmetal b. Polyvalent Binary Ionic – iii. Name – Roman Numerals! iv. Formula – Transition Metal & Nonmetal c. Polyatomic Ionic – v. Name – Metals & anions with “ate” or “ite” endings vi. Formula – Parentheses! Writing Formulas for Ionic Compounds Write the formula for the following ionic compounds: 70) Barium oxide BaO 71) Lithium nitride Li3N 72) Magnesium fluoride MgF2 73) Lithium phosphate Li3PO4 74) Calcium iodide CaI2 75) Copper (II) nitrite Cu(NO2)2 76) Beryllium chloride BeCl2 77) Gold (III) carbonate Au2(CO3)3 78) Potassium sulfate K2SO4 79) Calcium hydroxide Ca(OH)2 80) Barium chloride BaCl2 81) Lithium oxide Li2O 82) Magnesium oxide MgO 83) Calcium fluoride 84) Iron (II) chloride 85) Silver bromide CaF2 FeCl2 AgBr Naming Ionic Compounds Write the name for the following ionic compounds. Be sure to include parenthesis with Roman numeral for the transition metals: 86) AlBr3 87) Fe2(SO3)3 88) MgCl2 89) K3PO4 90) TiCl2 91) Cr2O3 92) Ag2S 93) Cu(OH)2 94) Zn3N2 95) AgNO3 96) PbF2 97) Cu2S 98) CaF2 99) FeCl3 100) Li3P 101) Au(C2H3O2)3 aluminum bromide iron (III) sulfite magnesium chloride potassium phosphate titanium (II) chloride chromium (III) oxide silver (I) sulfide copper (II) hydroxide zinc (II) nitride silver (I) nitrate lead (II) fluoride copper (I) sulfide calcium fluoride iron (III) chloride lithium phosphide Gold (III) acetate Review Covalent Bonding 102) How many valence electrons are in the following: a. Carbon 4 103) 104) 105) 106) b. Silicon 4 c. Chlorine 7 d. Nitrogen 5 What is the name given to the pairs of valence electrons that do not participate in bonding? Lone or unpaired What is the name given to the pairs of valence electrons that do participate in bonding? Shared pair In a covalent bond, electrons are shared between atoms. Which of the following is a covalent molecule (circle one)? a. LiCl b. BaO c. NO2 d. MgBr2 Writing Formulas for Covalent Compounds 107) Write the formula for the following covalent compounds: a. Sulfur tribromide SBr3 b. Dinitrogen tetrachloride N2Cl4 c. Dihydrogen monoxide H2 O d. Phosphorus Trihydride PH3 e. Tetracarbon octahydride C4H8 Nomenclature for Covalent Compounds 108) Name the following covalent compounds: a. N4O6 tetranitrogen hexoxide 109) b. SO3 sulfur trioxide c. SeF4 selenium tetrafluoride d. PCl5 phosphorous pentachloride e. XeI4 xenon tetraiodide How many electrons are shared in a single bond? Double bond? Triple bond? a. Single bond: 2 b. Double bond: 4 c. Triple bond: 6 110) Give 2 examples of a diatomic molecule. Hydrogen, oxygen, fluorine, nitrogen, chlorine, iodine, bromine 111) List 3 properties of covalent molecules. a. Low melting point b. 2 nonmetals c. Isn’t soluble 112) In a nonpolar covalent bond, electrons are shared equally. 113) In a polar covalent bond, electrons are shared unequally. 114) Label the partial positive and partial negative charge on the molecule below. 115) Define electronegativity (use the textbook). Ability of an atom to attract electrons 116) Electronegativity determines the type of bond. 117) Draw an example of a polar molecule. 118) Draw an example of a nonpolar molecule. 119) What is an intramolecular force? Provide 2 examples. The forces within a molecule are called intramolecular forces. This includes ionic and covalent bonds. 120) What is an intermolecular force? The forces between 2 molecules are called intermolecular forces. These forces are much weaker than intramolecular forces. 121) Label the intermolecular and intramolecular forces on the figure below: 122) Which force is the strongest, intermolecular or intramolecular? Intramolecular Why are polar molecules “sticky”? Because the partial charges are attracted to one another making them hard to separate. 123) 124) Fill in the chart below: Molecule: F2 Name: fluorine Polar or nonpolar bonds: nonpolar Lewis Structure: Molecule: HCl Name: hydrogen monochloride Polar or nonpolar bonds: polar Lewis structure: Molecule: CCl4 Name: carbon tetrachloride Polar or nonpolar bonds: polar Lewis structure: Molecule: SiO2 Name: silicon dioxide Polar or nonpolar bonds: polar Lewis structure: Molecule: HI Name: hydrogen monoiodide Polar or nonpolar bonds: polar Lewis structure: Molecule: N2H4 Name: dinitrogen tetrahydride Polar or nonpolar bonds: polar Lewis structure: Molecule: CO Name: carbon monoxide Polar or nonpolar bonds: polar Lewis structure: Molecule: CCl4 Name: carbon tetrachloride Polar or nonpolar bonds: polar Lewis structure: Molecule: H2O Name: dihydrogen monoxide Polar or nonpolar bonds: polar Lewis structure: Molecule: CO2 Name: carbon dioxide Polar or nonpolar bonds: polar Lewis structure: Unit 4 – Chemical Reaction 125) Which number represents the coefficient? 5 126) Which number represents the subscript? 2 127) How many atoms of Hydrogen are present? 10 128) What is a chemical reaction? A change that results in a substance with a different identity. 129) A substance that enters into a chemical reaction is called what? Reactants 130) A substance that is formed by a chemical reaction is called what? Products 131) T/F 132) What does the stand for? Yield 133) The law of conservation of mass states that mass can’t be created or destroyed. 134) Write the general equations for the 5 main types of chemical reactions. a. 5H2 In a true chemical reaction a new substance must form. True Synthesis = A + B AB b. Decomposition = AB A + B c. Single Replacement = A + BC AC + B d. Double Replacement = AB + CD AD + CB e. Complete Combustion = CxHy + O2 CO2 + H2O 135) Balance the following chemical reactions, label the reactants and products, and identify the type of reaction. a. b. c. d. e. f. g. EVERYTHING ON THE LEFT OF THE ARROW IS A REACTANT EVERYTHING ON THE RIGHT OF THE ARROW IS A PRODUCT 2Au2O3 → 4Au + 3O2 Reaction type: Decomposition 1SiC + 2Cl2 → 1SiCl4 + 1C Reaction type: Single 1S8 + 24F2 → 8SF6 Reaction type: Synthesis 3KOH + 1H3PO4 → 1K3PO4 + 3H2O Reaction type: Double 2H2 + 1O2 → 2H2O Reaction type: Synthesis 1S8 + 12O2 → 8SO3 Reaction type: Synthesis 1Zn + 2HCl → 1ZnCl2 + 1H2 Reaction type: Single 136) Which metals would be replaced in a solution, if a piece of lead metal were added? Cu, Hg, Ag, Au 137) You have a mixture that contains solutions of both Sodium and Aluminum compounds. You wish to ‘recover’ the aluminum but not the sodium. What other metal should you add to this mixture? Li, Ca, or Mg 138) Which of the metals could NOT be used to remove mercury from a nitrate solution? Ag or Au 139) Write the equation for “magnesium reacts with chlorine to produce magnesium chloride” Mg + Cl2 MgCl2 140) CuCl2 + Mg Cu + MgCl2 Which is the more active metal? Magnesium Unit 5 - Nuclear 141) On December 7, 1941, the Japanese attacked Pearl Harbor and the US declared war on Japan and then joined World War II. 142) What rare uranium isotope is useful in the fission reaction? U-235 143) Where did they locate the facility test site? 144) Who was selected to run the project from conception to completion? 145) The facility at Oak Ridge, TN was top secret. Why? Los Alamos Robert Oppenheimer Creation of Y-12 (separate U isotopes) – So Germans would not know about it 146) In 1941, what substance was created to be used to fuel the A-bomb? Plutonium 147) What president took over for Roosevelt and had to be updated on the secret? 148) In February 1945, the first bomb was completed which used U-235. What was it called? Little Truman Boy 149) What was the name given to the bomb that used plutonium? Fat Man 150) What code name was given to the first test of an atomic weapon? Trinity 151) At 5:30 am, on July 16, 1945, the world entered the “Atomic Age.” How powerful was the bomb? ~20,000 tons of TNT 152) On August 6, Enola Gay dropped Little Boy over Hiroshima. What percentage of the city was destroyed? 90%