Chemistry Fall Semester Review Key

Chemistry Fall Semester Review Key 1.) Write the number of significant figures that are in each of these correctly taken measurements. 1. _____2___ 2300 7. ___3_____ 0.0230 2. ___3______ 8.10 x 10 m 8. ___4______ 800.0 3. ____2_____ 5.0 x 10 9. ___4______ 0.0003400 7 -2 4. _____4____ 1003 10. __5_______ 990.00 5. _____4____ 2.110 x 10 11. __2_______ 450 6. _____5____ 8.0335 12. __4_______ 6.100 x 10 -3 -2 2.) Add, subtract, multiply, or divide the following using significant figure rules. a) 258.3 kg + 257.11 kg + 253 kg = 768 12.12 f) 93.26 cm – 81.140 cm = b) 13.65 cm x 23.20 cm x 42.0 cm = 13300 g) 4.84 g / 2.4 mL = 2.0 c) 120. m x 0.10 m = 12 3.00 s h) 60.2 m / 20.1 s = d) 5.46 g + 1.329 g – 3.12 g = 3.67 cm = 1.1 cm i) 2.53 cm – 1.4 e) 0.560.030= 19 7.91) x 5.3 = 89 j) (8.807 + 3.) Write the following quantities from scientific notation to ordinary (standard) notation with the correct number of significant figures. a. b. c. d. 5.80100 x 10 s 580,100 s 8.12 x 10 m 81,200,000 m 5.0 x 10 kg 0.050 kg 2.110 x 10 m 0.002110 m 5 7 -2 -3 4.) Express the following quantities in scientific notation with the correct number of significant figures. a. 5,800 m 5.8x10 m b. 450,000 m 4.5x10 m c. 302,000,000 m 3.02x10 m d. 86,100,000,000 m 8.61x10 m e. 0.000508 kg 5.08x10 kg f. 0.000000450 kg 4.50x10 kg g. 0.0003600 kg 3.600x10 kg h. 0.004 kg 4x10 kg 3 5 8 10 -4 -7 -4 -3 5.) Use Dimensional Analysis to solve the following problems. Show all work. Remember to use sig figs and box your final answer. a) Convert 360 centimeters to nanometers. 360 cm x 10-2 m1 cm x 1 nm10-9 mn=3.6 x 109 b) Convert 600.0 megahertz to Hertz. 600.0 MHz x 106 Hz1 MHz =6.00 x 108 c) Convert 3.4 mm to km. 3.4 mm x 10-3 m1 mm x 1 km103 mn=3.4 x 10-6 d) A measurement of 1.40 x 10 Hz is equivalent to how many MHz (megahertz)? 1.40 x 106 Hz x 1 MHz106 Hz = 1.40 6 e) How many meters are in 4.7 x 10 nm? 4.7 x 10-7 nm x 10-9 m1 nm =4.7 x 10-16 -7 f) Convert 672 micrograms to milligrams. 672 g x 10-6 g1 g x 1 mg10-3g =6.72 x 102 g) How many nanograms are in 7.3 x 10 μg ? 7.3 x 106 g x 10-6 g1 g x 1 ng10-9g =7.3 x 109 6 6.) Give the measurements for each of the following using the correct number of significant figures (include units): a.) 3.0 cm b.) 4.70 cm 7.) Compare and contrast accuracy and precision and draw and example of each. When scientists look at measurements, they want to know how accurate and how precise the measurements are. Accuracy is how close a measured value is to an accepted value. Precision is how close a series of measurements are to one another. Precise measurements might not be accurate, and accurate measurements might not be precise. When you make measurements you want to aim for both accuracy and precision. 8.) Complete the following table for phases of matter. States (Phases) of Matter Phase Gas Shape Volume indefinite indefinite Particle Spacing/ Compressibility Particle Movement spread out; compressible rapid Diagram of Particles Liquid indefinite Solid definite definite compact; incompressible less than gas more than solid definite compact; incompressible slow 9.) For each of the following statements, indicate whether it refers to a solid, liquid, or gas. a. Solid b. Gas c. Liquid d. Gas e. Liquid f. Solid Particles vibrate in place. Particles are moving past each other at high speeds Particles are close together but able to move past each other This state of matter is the most compressible This state of matter has a definite volume but an indefinite shape. This state of matter is the most dense. 10.) Density of Matter. If you consider the solid, liquid, and gas state of one particular substance, this rule holds true in most cases. Draw an example of the particle spacing in each. Solid Liquid Gas Most dense dense Least 11.) Draw the flowchart for the types of matter. Include the terms matter, pure substances, mixtures, elements, compounds, homogeneous mixture, and heterogeneous mixtures. 12.) Label each model below as an element, compound, or mixture. Mixture Compound Element Mixture 13.) Types of Matter a.) Explain the difference between an element and a compound. Give two examples of each. Compounds are two or more elements which are chemically combined. Element: Oxygen, Chlorine Compound: CO , H O 2 2 b.) Explain is the difference between a heterogeneous mixture and a homogeneous mixture and the separation techniques that can use for each. Give two example of each type of mixture. A heterogeneous mixture is one that is not blended smoothly throughout. Examples of heterogeneous mixtures include smoky air and muddy water. A homogenous mixture is one that has a constant composition throughout. By dissolving sugar in water, you create a homogenous mixture. 14.) Complete the following table: Term Intensive Property Extensive Property Physical Property Chemical Property Description A property which does not change with size A property which changes with size A property which can be measured without altering the material. A property which describes how the material will or will not react with other materials. A change which changes only a physical property of the material Physical Change Chemical Change A change which changes the actual material 15.) Complete the following table: Property Extensive or Intensive Physical or Chemical Color Intensive Combustibility Intensive Hardness Intensive Physical Density Intensive Physical Mass Physical Chemical Extensive Physical Melting Point Intensive Physical Ductility Intensive Physical Volume Extensive Reactivity with Acid Intensive Odor Intensive Weight Physical Chemical Physical Extensive Malleability Intensive Tendency to Corrode Intensive Physical Physical Chemical 16.) Describe the five evidences of a chemical change. Light is created Heat is absorbed/evolved Gas is created • smoke, bubbles, fizzing, new odor, etc Color change occurs (unexpected, i.e. yellow + blue makes black) Precipitate Electrical signal generated 17.) Classify the following changes of matter as Physical or Chemical and provide evidence for each. Scenario An ice cube is melted Sugar is dissolved in water Baking soda reacts with vinegar and forms a gas A piece of copper wire is bent in half An aspirin is crushed into fine powder Copper turns green when exposed to the environment Physical or Chemical Physical Evidence Physical Chemical It’s still H O (same chemical make-up) It’s still sugar and H O Bubbles, gas Physical It’s still Copper Physical It is still aspirin Chemical Color change (coppery orange to green) 2 2 Two clear liquids are mixed and a yellow, cloudy substance forms A tree burns to form ashes A piece of paper is crumpled up Water freezes to form ice A candle burns A candle melts A car rusts A cup of water evaporates Chemical Precipitate Chemical new substance - ash; color change (black), heat evolved It’s still paper It’s still H O heat evolved; light emitted, gas (smoke) It’s still wax color change (rusty -red); new substance with different physical properties change of state (liquid → gas) Physical Physical Chemical Physical Chemical Physical 2 18.) Complete the following table for Atomic Theory. Scientist John Dalton J.J. Thomson Description of Experiment/Contributions 5 Postulates, 1. atoms are indivisible 2. atoms are indestructible 3. all atoms of an element are the same 4. atoms combine in whole number ratios to make compounds 5. A chemical reaction is a rearrangement of atoms “Father of Atomic Theory” Cathode Ray Tube demonstrated that cathode rays could be deflected, or bent, by magnetic or electric fields, which indicated that cathode rays consist of charged particles. More important, by measuring the extent of the deflection of the cathode rays in magnetic or electric fields of various strengths, Thomson was able to calculate the massto-charge ratio of the particles. These particles were emitted by the negatively charged cathode and repelled by the negative terminal of an electric field. Because like charges repel each other and opposite charges Models Atomic model and theory, Dalton was the first person to recognize a workable distinction between the fundamental particle of an element (atom) and that of a compound (molecule), “Billiard Ball Model” (atom tiny, hard sphere) Discovered the electron, Plum Pudding model attract, Thomson concluded that the particles had a net negative charge; these particles are now called electrons. Ernest Rutherford Niels Bohr Gold Foil Experiment - aimed a stream of α particles at a very thin gold foil target and examined how the α particles were scattered by the foil. A small fraction of the α particles were deflected at large angles, and some were reflected directly back at the source.It suggested that the mass of the gold atom is concentrated in a very small region of space, which he called the nucleus. He proposed a model of the atom that is similar to the model of the solar system. The electrons go around the nucleus like planets orbit around the sun. Electrons behave like waves and have a certain energy. All electrons have their energy levels – a certain distance from the nucleus. Each energy level can hold a certain number of electrons. Discovered 1.) atom mostly made up of empty space, 2.) small, dense, positive nucleus Pioneered the quantum model; Bohr Model This model can be portrayed as a nucleus surrounded by an electron cloud. Electrons are NOT in circular orbits around nucleus. Quantum Mechanical -Electrons are in a 3-D region Model around the nucleus called atomic orbitals. -The atomic orbital describes the probable location of the electron 19.) Average Atomic Mass Practice a. Zinc has 5 isotopes that occur in nature as shown in the table. Find the average atomic mass. Isotope Mass of Isotope (amu) Percent abundance (%) Zinc-64 63.929 48.89 Zinc-66 65.927 27.81 Zinc-67 66.927 4.11 Zinc-68 67.925 18.57 Zinc-70 69.925 0.62 63.929 (.4889) + 65.927 (.2781) + 66.927 (.0411) + 67.925 (.1857) + 69.925 (.0062) = 65.39 amu b. What is the average atomic mass of element X given the following abundance information on the isotopes of X? Identify the element. Mass number Abundance X-28 92.21% X-29 4.70% X-30 3.09% 28(.9221) + 29(.0470) + 30(.0309) = 28.1 amu Silicon 20.) Complete the element, symbol, atomic number, the average atomic mass, mass number, protons, electrons, and neutrons of the most common isotope for each of the following elements. Hyphen Isotope Notation Symboli c Isotopic Notatio n Atomi c# Uranium -235 92235U 92 Atomic Mass (periodi c table) Mas s# 238.03 235 Proton s Electro ns Neutro ns (mass #-p ) Neutra l, Cation, or Anion 143 Neutra l + 92 92 Oxygen18 818O 8 15.99 18 8 8 10 Neutra l Bismuth213 83213B i 83 208.98 213 83 86 130 Anion Vanadium50 2350V 23 50.94 50 23 23 27 Neutra l Magnesiu m-25 1225Mg 2+ 12 24.31 25 12 10 13 Cation Krypton85 3685Kr 36 83.80 85 36 36 49 Neutra l Gold 79198A u 79 196.97 198 79 79 119 Neutra l Thorium 90229T h 90 232.04 229 90 90 139 Neutra l 3- 21.) Give the “p-e-n” numbers for each of the following ions: a.) Zn b.) Ag 2+ 1+ p- 30 e- 28 n- 35 c.) P d.) I 3- p- 15 e- 18 n- 16 p- 47 e- 46 n- 61 1- p- 53 e- 54 n- 74 22.) Complete the following table for each types of radiation. Radiation Symbol(s) Charge 24⍺ +2 alpha Mass Penetrates Energy 4 amu Low Penetration Low Energy -10β -1 Negligible Medium Penetration Medium Energy Ɣ No Charge No Mass High Penetration High Energy beta gamma Illustration 23.) Compare and contrast fission and fusion. Include current application of each. 24.) Complete the following reactions and identify the type of decay. 86222Rn → 24He + 84218Po Type of decay: Alpha 82214Pb → 83214Bi +-10e Type of decay: Beta 1427Si 83214Bi → -10e + 1527P → 24He+ 81210Tl 92234U → 93234Np + -10e 2760Ca → 2760Ca + 00 Ɣ Type of decay: Beta Type of decay: Alpha Type of decay: Beta Type of decay: Gamma 25.) Write a balanced nuclear equation for the Beta decay of the following: a. Carbon-14 614C → 714N + -10e b. Strontium-90 3890Sr → 3990Y + -10e c. Potassium-40 1940K → 2040Ca + -10e 26.) Write a balanced nuclear equation for the Alpha decay of the following: a. Thorium-234 90223Th → 24He+ 88219Ra b. Radium-222 88222Ra → 24He+ 86218Rn c. Polonium-214 84214Po → 24He+ 82210Pb 27.) How are the mass and atomic number of the nucleus affected by the each the following? a. Alpha particle Mass # reduced by 4, Atomic # reduced by 2. b. Beta particle Mass number is unchanged, Atomic # is increased by 1 28.) Draw and explain Aufbau’s Principle. Orbitals are filled in order of increasing energy; electrons will fill lowest energetically available orbitals first n: 1 < 2 < 3 <4 l: s < p < d < f m: p = p = p l x y z 29.) Complete the following table using the Periodic Table. Noble Gas Electron Configuration Group Period [Ne] 3s 3p 2 3 [Ar] 4s 3d 4p 2 10 [Ar]4s 2 [Kr]5s 4d 5p 2 4 10 5 Element 5A 3 Phosphorus 6A 4 Selenium 2A 4 7A 5 Calcium Iodine 30.) Write the extended electron configurations for each element: a.) Sodium 1s 2s 2p 3s 2 2 6 1 b.) Chlorine 1s 2s 2p 3s 3p 2 2 c.) Neon 1s 2s 2p 2 2 6 6 2 5 31.) Draw the Lewis Dot Structure for each of the following elements and identify each element’s family/group: a.) Potassium Alkali Metal b.) Fluorine Halogen c.) Argon Noble Gas d.) Magnesium Alkaline Earth Metal 32.) Describe the properties of the following families in the periodic table: a.) Alkali Metals- The alkali metals have 1 valence electron. They are soft metallic solids, with characteristic metal properties of shiny luster and high thermal and electrical conductivities. The alkali metals have low densities and melting points, increasing and decreasing respectively with increasing atomic mass. The alkali metals react vigorously with water, producing hydrogen gas and a solution of an alkali metal hydroxide. These reactions are very exothermic, in many cases generating enough heat to ignite the H , producing a fire or sometimes even an explosion. Sodium and potassium are abundant in all biological systems. Hydrogen is in the Alkali Metals Family - but; Although hydrogen has 1 valence electron and is found on the periodic table in Group 1A, it does not truly belong to any particular group. Because it has no core of electrons producing nuclear shielding, its ionization energy is markedly higher than that of the alkali metals. Where alkali metals lose their valence electron, hydrogen normally shares its valence electron. The fact that hydrogen can actually gain an electron when forming hydrides further illustrates that it is not truly a member of the alkali metal family 2 b.) Alkaline Earth Metals- The alkaline earth metals have 2 valence electrons. They have characteristic metal properties but are harder, more dense, and melt at higher temperatures than the alkali metals. Their reaction with water increases down the group. Beryllium does not react with water, even when heated red-hot. Magnesium does not react with liquid water but does react with steam. Calcium and the elements below it react with water at room temperature, although more slowly than the alkali metal adjacent to them on the periodic table. Magnesium and calcium are essential for living organisms. 99% of the calcium in the human body is found in the skeletal system. c.) Halogens- The halogens have 7 valence electrons. They are typical nonmetals. Their melting and boiling points increase with increasing atomic number. Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid. Halogens are highly active with high electron affinities. In fact, the letter X is sometimes used in chemical equations to indicate any one of the halogen elements. Fluorine gas is so reactive that it is difficult and dangerous to use, requiring specialized lab equipment. Unlike fluorine, chlorine reacts slowly with water to form relatively stable aqueous solutions. Chlorine is often added to drinking water and swimming pools, where the HOCl that is produced serves as a disinfectant. Halides are all very soluble in water and dissolve to form the hydrohalic acids (aq) d.) Noble Gases- The noble gases have 8 valence electrons. Because the noble gases possess such stable electron configurations, they are exceptionally unreactive. In fact, until the early 1960s the elements were called the inert gases because they were thought to be incapable of forming chemical compounds. Xenon is now known to react with F to form the molecular compounds XeF , XeF , XeF . Only one stable compound of krypton is known, KrF . 2 (g) 2 4 6 2 33.) Elements in the same group/family have the same physical and chemical properties. Why? All elements in a representative group have the same number of valence electrons. 34.) Compare and contrast Mendeleev’s and Moseley's arrangement of the Periodic Table. Mendeleev arranged the periodic table periodically by atomic mass. We now know that this way does not provide the best results. He is important because he is the first one to leave gaps in the table and to predict properties of undiscovered elements. He predicted the properties of Gallium before the element was discovered. Henry Moseley developed the concept of atomic numbers - correctly stating that the atomic number was equal to the number of protons in the nucleus of an atom and the number of electrons surrounding the nucleus. Henry Moseley determined the atomic number of each of the element. He modified the Periodic Law to read that the properties of the elements vary periodically with their atomic numbers. Moseley's modified Periodic Law puts the elements tellurium and iodine in the right order, as it does for argon and potassium, cobalt and nickel 35.) Answer the Electromagnetic Radiation (EMR) questions below: a.) List all the forms of EMR including all the colors of visible light from: 1.) Lowest frequency to highest frequency Radio Micro Infra Visible UV X-Ray Gamma 2.) Shortest wavelength to longest wavelength G X UV V IR M R b.) Describe the relationship between energy and frequency? Directly Proportional c.) Describe the relationship between wavelength and frequency? Inversely Proportional 36.) Use the constants below to solve the following Wave Calculations: a.) Calculate the frequency of a wave which has a wavelength of 4.15 meters. c=f 3.0*10 =(4.15)(f) f=3.0*1084.15=7.22 x 10 8 7 b.) What is the wavelength of a wave with a frequency of 5.77 x 10 Hz? 14 c = λf λ =cf λ =3.00 x 108 m/s5.77 x 1014 Hz λ = 5.20 x 10 m -7 c.) What is the energy of a photon with a wavelength of 7.67 x 10 ? E = h x cλ E = 6.63 x10-34 Jᐧs x 3.00 x 108m/s7.67 x 10-7 m E = 2.59 x 10 J -7 -19 d.) An X-ray photon has an energy of 2.99 x 10 J. What is its wavelength? -18 E = hf f = Eh f = 2.99 x 10-18 J6.63 x10-34 Jᐧs f = 4.01 x 10 Hz 15 E = h x cλ λ =h cE λ =6.63 x10-34 Jᐧs x 3.00 x 108m/s2.99 x 10-18 J λ = 6.62 x 10 m -8 e.) What is the wavelength of a wave that gives off 5.6 x 10 Joules of energy? E = h x cλ E = 6.63 x10-34 Jᐧs x 3.00 x 108m/s5.6 x 1023 m λ = 3.6 x 10 m 23 -49 37.) Complete the following table on metals, nonmetals, and metalloids. Classification Location on Periodic Table Properties Examples Metal Nonmetal Metalloid Many, but not all, Periodic Tables have a line dividing these two. The dividing line begins between boron and aluminum and stair-steps down and to the right, one square at a time. Metals are left of the line. Many, but not all, Periodic Tables have a line dividing these two. The dividing line begins between boron and aluminum and stair-steps down and to the right, one square at a time. Nonmetals are to the right. When moving from left to right across the Periodic Table, metal characteristics decrease and nonmetal characteristics increase. In the area of the dividing line, some elements have both metallic and nonmetallic properties. Most metallic elements have a shiny luster, have high melting points, are good conductors of heat and electricity, are malleable (can be pounded into thin sheets), are ductile (can be drawn into wires). Metals tend to have low ionization energies and lose electrons relatively easily to form positive ions. Nonmetals vary greatly in appearance. Most are not lustrous, have low melting points, are poor conductors of heat and electricity, and not malleable or ductile. Because of their electron affinities, nonmetals tend to gain electrons relatively easily to form negative ions. Metalloids touch the metal/nonmetal line on the Periodic Table and have properties between those of metals and nonmetals. They may have some characteristic metallic properties but lack others. Strontium, Gold, Osmium Carbon, Chlorine, Fluorine Several metalloids, most notably silicon and germanium, are electrical semiconductors and the principal elements used to manufacture integrated circuits and computer chips. 38.) Complete the following table on periodic trends. Trend What is it? Trend across a period Trend across a group What causes trend? a measure of the size of an atom decrease in size increase in size Atomic Radius Ionization Energy Energy is required to pull an electron away from an atom. Ionization energy increases across a period Ionization energy decreases as you go down a group As you go down the table, there are more energy levels (rings). The valence electrons are located further from the nucleus. All atoms on a row of the periodic table have their valence electrons on the same energy level. In other words, they are the same distance from the nucleus. However, each element adds another proton to the nucleus. The inner core does not change. Therefore, the effective nuclear charge increases. According to Coulomb’s Law, this increase in charge will increase the force. Energy is required to pull an electron away from an atom. The first ionization energy of an element is the amount of energy required to pull the first valence electron away from an atom of the element. Atoms with high ionization energies, such as fluorine, oxygen, and chlorine, are found on the right side of the periodic table and are unlikely to form positive ions by losing electrons. Instead, they usually gain electrons, forming negative ions. Atoms with low ionization energies, such as sodium, potassium, and strontium, lose electrons easily to form positive ions and are on the left side of the periodic table. Recall that atoms decrease in size from left to right across a period. First ionization energies generally increase across a period of elements primarily because the electrons to be removed are successively closer to the nucleus. First ionization energies decrease moving down through a group of elements because the sizes of the atoms increase and the electrons to be removed are farther from the nucleus. The tendenc y for an atom to attract bonding electron s Electronegativit y Electronegativit y increases as you go across a period Electronegativit y decreases as you go down a group When atoms combine chemically with each other, they do so by forming a chemical bond. This bond involves either the transfer of electrons or sharing of electrons to varying degrees. The nature of the bond between two atoms depends on the relative ability of each atom to attract electrons from the other, a property known as electronegativity . The maximum electronegativity value is 3.98 for fluorine, the element that attracts electrons most strongly in a chemical bond. The trends in electronegativity in the periodic table are generally similar to the trends in ionization energy. The lowest electronegativity values occur among the elements in the lower left of the periodic table. These atoms, such as cesium, rubidium, and barium, are large and have few valence electrons, which they lose easily. Therefore, they have little attraction for electrons when forming a bond. Elements with the highest electronegativity values, such as fluorine, chlorine, and oxygen, are found in the upper right of the periodic table (excluding, of course, the noble gases, which do not normally form chemical bonds). These atoms are small and can gain only one or two electrons to have a stable noble-gas configuration. Therefore, when these elements form a chemical bond, their attraction for electrons is large. Electronegativity generally increases across a period and decrease down through a group. 39.) List the order of the atoms using the given trends (indicate highest and lowest): Trend Elements List Order Largest Atomic Radius (Lowest to Highest) Lowest Ionization Energy (Lowest to Highest) Least Electronegative (Lowest to Highest) Highest Ionization Energy (Lowest to Highest) Most Electronegative (Lowest to Highest) Smallest Atomic Radius (Lowest to Highest) Ca, K, Br Mg, Sr, Ba C, Li,O I, Cl, F, P, Si, S Br, I, Cl Br, Kr, Se Al, Li, Cl As, F, Se Br, Ca, K Mg, Sr,Ba Li, C, O I, Cl, F Si, P, S I, Br, Cl Se, Br, Kr Li, C, F Fr, Fe, F Li, Al, Cl F, Se, As Sn, I, Sr I, Sn, Sr C, Li, F Fe, F, Fr 40.) Complete the following table: Eleme nt symbol Cesium Catio n or Anio n C Grou p# Valence Electro ns Oxidatio n Number Ion symb ol Cs 1A 1 1+ Cs Cesium Strontium Sr 2A 2 2+ Sr C Bromine Br 7A 7 1- Br Strontiu m Bromide Boron B 3A 3 3+ B Boron C Phosphoro us Oxygen P 5A 5 3- P A O 6A 6 2- O Phosphi de Oxide Element Name +1 +2 -1 +3 -3 -2 Ion Name A A Iodine I 7A 7 1- I Lithium Li 1A 1 1+ Li -1 +1 Iodide A Lithium C 41.) Complete the following vocabulary descriptions and terms. Term Description Example Positively charged ion consisting of one atom, usually a metal Ca , Na Negatively charged ion consisting of one atom, usually a nonmetal C,S ion consisting of two or more atoms PO , NO The charge of the ion 3+, 2-, 1+ Metals that have more than one oxidation number, Roman Numerals are used to indicate charge Iron (III) Fe , Copper(I) Cu +2 +1 Monatomic Cation Monatomic Anion -- 4 -3 -2 3 - Polyatomic Ion Oxidation Number Transition Metal 3+ +1 42.) Describe properties of ionic compounds. Crystal Lattice structure allows for properties of ionic compounds including:high melting point, high boiling point, nonconductors as solids, soluble in water, dissociate into ions known as electrolytes, hard, and brittle 43.) For the following combinations, draw the electron dot diagrams (Lewis structures) and show the transfer of electrons. Also, write the chemical formula and name. a. aluminum and oxygen b. calcium and chlorine c. sodium and sulfur d. lead(III) and nitrogen 44.) Use the ΔEN to predict whether the following bonds are ionic or covalent. Remember, if ΔEN is greater than 1.7, we consider it to be ionic. a. Al S d. SO 2.5 - 1.5 = 1 = covalent 3.5 - 2.5 = 1 = covalent b. NO e. KCl 3.5 - 3.0 = .5 = covalent 3.0 - 0.8 = 2.2 = ionic c. CaF f. CH 2 3 2 2 4 4.0 - 1.0 = 3 = ionic 2.5 - 2.1 = .4 = covalent 45.) Write the ions and chemical formulas for the following ionic compounds: a. barium chloride BaCl b. 2 potassium nitride KN 3 c. MgI magnesium iodide 2 d. aluminum phosphide AlP e. cesium fluoride CsF f. calcium oxide CaO 46.) Name the following ionic compounds: a. MgCl Magnesium chloride b. Ca P Calcium phosphide c. Al O Aluminum oxide d. NaF Sodium fluoride 2 3 2 2 3 e. KS Potassium sulfide f. Cs N Cesium nitride 2 3 47.) Write the ions and chemical formulas for the following ionic compounds: a. manganese (VII) oxide Mn O +7 -2 Mn O 2 7 b. copper (II) chloride Cu Cl CuCl +2 - 2 c. titanium (III) phosphide Ti P TiP +3 -3 d. iron (III) sulfide Fe S +3 -2 Fe S 2 3 e. cobalt (II) bromide Co Br +2 - CoBr 2 f. lead (IV) oxide Pb O +4 -2 Pb O 2 4 PbO 2 48.) Name the following ionic compounds: a. Mn S Manganese (VII) sulfide b. CuF Copper (II) fluoride 2 7 2 c. Fe O 3 Iron (III) oxide d. Zn P 2 Zinc phosphide e. PbCl 4 Lead (IV) chloride 2 3 f. CrO Chromium (VI) oxide g. NiO Nickel (II) oxide h. Ag S Silver sulfide 3 2 49.) Write the ions and chemical formulas for the following ionic compounds: a. sodium chromate Na CrO + 4 -2 Na CrO 2 4 b. barium nitrate Ba NO +2 3 Ba(NO ) 3 - 2 c. ammonium sulfate NH 4 + SO 4 -2 (NH ) SO 4 2 4 d. aluminum hydroxide Al +3 OH - Al(OH) 3 e. calcium phosphate Ca PO +2 4 Ca (PO ) 3 4 -3 2 f. copper(II) sulfate Cu SO +2 4 -2 CuSO 4 g. iron(II) nitrite Fe NO +2 2 Fe(NO ) 2 2 h. tin(IV) acetate Sn CHO +4 2 3 2 - Sn(C H O ) 2 3 2 4 50.) Name the following ionic compounds: a. NH Cl Ammonium chloride 4 b. K Cr O Potassium dichromate c. BeSO Beryllium sulfate 2 2 7 4 d. (NH ) N e. NaClO Ammonium nitride Sodium chlorate f. Cu SO Copper (I) sulfate 4 3 3 2 4 g. Pb (PO ) Lead (IV) phosphate 3 4 4 h. MgSO 3 Magnesium sulfite 51.) Complete the following table: Cation Anion Chemical Formula Chemical Name Ti 5+ N Ti N 3- 3 Titanium (V) nitride 5 Ca +2 NO - Ca(NO ) Ni +2 CO -2 NiCO Al OH H 2+ + 4 Mg + Hg + Calcium nitrate 2 Nickel (II) Carbonate 3 Al(OH) - Aluminum hydroxide 3 - 3 Sr(HCO ) 3 2 Strontium bicarbonate PO 3- H PO SO -2 (NH ) SO Ammonium sulfate MgO Magnesium oxide 4 3 Hydrogen phosphite 3 4 2 O +2 +3 3 HCO 3 NH Al 3 +3 Sr 3 -2 4 CHO Al(C H O ) Aluminum acetate NO HgNO Mercury(I) nitrate KP potassium phosphide 2 2 3 3 1- 2 3 2 3 3 K 1+ P Al +3 NO Al(NO ) Aluminum Nitrite + N Ag N Silver nitride +2 PO Ag Ca 3- 2 3- 4 Rb + O Pb 4+ OH F Ca SO +2 -3 1- -2 + Br Mg +2 ClO 2+ Cl Mg Cu 3 4 2 - 4 - 1- Calcium Phosphate Rubidium oxide 2 Na Be Ca (PO ) Rb O - 3 2 3 3 -2 Sn +2 3 Pb(OH) Lead (IV) hydroxide SnF Tin (II) fluoride 4 2 Ca(SO ) Calcium Sulfite NaBr Sodium bromide 3 Mg(ClO ) Magnesium Perchlorate BeCl Beryllium chloride 4 2 2 +2 NO Mg(NO ) Magnesium nitrite +2 Br CuBr Copper (II) Bromide 2 - 2 2 2 Ag CrO + 4 Pb 2+ S Ca 2+ PO +3 CrO 2- 4 Ag CrO Silver chromate PbS Lead (II) sulfide 2 4 Ca (PO ) 3- 4 Fe -2 3 -2 4 2 Fe (CrO ) 2 4 3 Iron (III) Chromate Na CO Al CHO Al(C H O ) Aluminum Acetate NO Zn (NO ) Zinc nitrate + 3 +3 Zn 2 2+ Na CO Calcium phosphate -2 2 2 3 3 1- 2 Sodium carbonate 3 3 2 3 3 2 52.) Give the oxidation number for each element in the following substances. a). Cl 2 Cl = 0 b). N O 2 3 N = +3 O = -2 c.) Mg (AsO ) 3 4 2 Mg = +2 As = +5 O = -2 c). H SO 2 3 H = +1 S = +4 O = -2 d). C C=0 e). Al S 2 3 Al = +3 S = -2 f). Fe(ClO ) 2 3 Fe = +3 Cl = +3 O = -2 g). MnO 2 Mn = +4 O = -2 h). KClO K = +1 Cl = +1 O = -2 i). CS 2 C = +4 S = -2 j). Zn(ClO ) 3 2 Zn = +2 Cl = +5 O = -2 k). Ni(NO ) 2 2 Ni = +2 N = +3 O = -2