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Chapter 3: Elements, Compounds, and the Periodic Table

Chapter 3: Elements, Compounds, and the Periodic Table. Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop. Discovery of Subatomic Particles. Late 1800s and early 1900s Cathode ray tube experiments showed that atoms are made up of subatomic particles

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Chapter 3: Elements, Compounds, and the Periodic Table

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  1. Chapter 3: Elements, Compounds,and the Periodic Table Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop

  2. Discovery of Subatomic Particles • Late 1800s and early 1900s • Cathode ray tube experiments showed that atoms are made up of subatomic particles • Discovered negatively charged particles moving from the cathode to the anode • Cathode – negative electrode • Anode – positive electrode

  3. Discovery of Electron JJ Thomson (1897) • Modified cathode ray tube • Made quantitative measurements on cathode rays • Discovered negatively charged particles • Electrons (e–) • Determined charge to mass ratio (e/m) of these particles • e/m = –1.76 x 108 coulombs/gram

  4. Millikan Oil Drop Experiment • Determining charge on Electron • Calculated charge on electron • e = –1.60 × 10–19 Coulombs • Combined with Thomson’s experiment to get mass of electron • m = 9.09 × 10–28 g

  5. Discovery of Atomic Nucleus Rutherford’s Alpha Scattering Experiment • Most alpha () rays passed right through gold • A few were deflected off at an angle • 1 in 8000 bounced back towards alpha ray source • Gave us current model of nuclear atom

  6. Discovery of Proton • Discovered in 1918 in Ernest Rutherford’s lab • Detected using Mass Spectrometer • Hydrogen had mass 1800 times the electron mass • Masses of other gases whole number multiples of mass of hydrogen Proton • Smallest positively charged particle

  7. Rutherford’s Nuclear Atom • Demonstrated that nucleus: • has almost all of mass in atom • has all of positive charge • is located in very small volume at center of atom • Very tiny, extremely dense core of atom • Where protons (p) and neutrons(1n) are located

  8. Atomic Structure • Electrons (e–) • Very low mass • Occupy most of atom’s space • Balance of attractive and repulsive forces controls atom size • Attraction between protons (p) and electrons (e–)holds electrons around nucleus • Repulsion between electrons helps them spread out over volume of atom • In neutral atom • Number of electrons must equal number of protons • Diameter of atom ~10,000 × diameter of nucleus

  9. Discovery of Neutron • First postulated by Rutherford and coworkers • Estimated number of positive charges on nucleus based on experimental data • Nuclear mass based on this number of protons always far short of actual mass • About ½ actual mass • Therefore, must be another type of particle • Has mass about same as proton • Electrically neutral • Discovered in 1932 by Chadwick

  10. Properties of Subatomic Particles • Three kinds of subatomic particles of principal interest to chemists Nucleus (protons + neutrons) Electrons

  11. Atomic Notation Atomic number (Z) • Number of protons that atom has in nucleus • Unique to each type of element • Element is substance whose atoms all contain identical number of protons • Z= number of protons Isotopes • Atoms of same element with different masses • Same number of protons ( ) • Different number of neutrons ( )

  12. Atomic Notation Isotope Mass number (A) • A = (number of protons) + (number of neutrons) • A = Z + N • For charge neutrality, number of electrons and protons must be equal Atomic Symbols • Summarize information about subatomic particles • Every isotope defined by two numbers Z and A • Symbolized by Ex.What is the atomic symbol for helium? He has 2 e–, 2 n and 2 pZ = 2, A = 4

  13. Isotopes • Most elements are mixtures of two or more stable isotopes • Each isotope has slightly different mass • Chemically, isotopes have virtually identical chemical properties • Relative proportions of different isotopes are essentially constant • Isotopes distinguished by mass number (A): e.g. • Three isotopes of hydrogen (H) • Four isotopes of iron (Fe)

  14. Example: What is the isotopic symbol for Uranium-235? • Number of protons (p) = 92 = number of electrons in neutral atom • Number of neutrons (1n) = 143 • Atomic number (Z) = 92 • Mass number (A) = 92 + 143 = 235 • Chemical symbol = U • Summary for uranium-235:

  15. Your Turn! An atom of has ___ protons, ___ neutrons, and ___ electrons. • 82, 206, 124 • 124, 206, 124 • 124, 124, 124 • 82, 124, 82 • 82, 124, 124

  16. Learning Check: • Fill in the blanks: symbol neutrons protons electrons 60Co 81Br 36 29 29 33 27 27 46 35 35

  17. Carbon-12 Atomic Mass Scale • Need uniform mass scale for atoms Atomic mass units(symbol u) • Based on carbon: • 1 atom of carbon-12 = 12 u (exactly) • 1 u = 1/12 mass 1 atom of carbon-12 (exactly) Why was 12C selected? • Common • Most abundant isotope of carbon • All atomic masses of all other elements ~ whole numbers • Lightest element, H, has mass ~1 u

  18. Calculating Atomic Mass • Generally, elements are mixtures of isotopes e.g. Hydrogen Isotope Mass % Abundance 1H 1.007825 u 99.985 2H 2.0140 u 0.015 How do we define atomic mass? • Average of masses of all stable isotopes of given element How do we calculate average atomic mass? • Weighted average • Use isotopic abundances and isotopic masses

  19. Learning Check Naturally occurring magnesium is a mixture of 3 isotopes; 78.99% of the atoms are 24Mg (atomic mass, 23.9850 u), 10.00% of 25Mg (atomic mass, 24.9858 u), and 11.01% of 26Mg (atomic mass, 25.9826 u). From these data calculate the average atomic mass of magnesium. 0.7899 x 23.9850 u = 18.946 u 24Mg 0.1000 x 24.9858 u = 2.4986 u 25Mg 0.1101 x 25.9826 u = 2.8607 u26Mg Total mass of average atom = 24.3053 u rounds up to 24.31 u

  20. Your Turn! A naturally occurring element consists of two isotopes. The data on the isotopes: isotope #1 68.5257 u 60.226% isotope #2 70.9429 u 39.774% Calculate the average atomic mass of this element. • 70.943 u • 69.487 u • 69.526 u • 69.981u • 69.734 u 0.60226 × 68.5257 u = 41.270 u 0.39774 × 70.9429 u = 28.217 u 69.487 u

  21. Periodic Table • Summarizes periodic properties of elements Early Versions of Periodic Tables • Arranged by increasing atomic mass • Mendeleev (Russian) and Meyer (German) in 1869 • Noted repeating (periodic) properties Modern Periodic Table • Arranged by increasing atomic number (Z): • Rows called periods • Columns called groups or families • Identified by numbers • 1 – 18 standard international • 1A – 8A longer columns and 1B – 8B shorter columns

  22. Modern Periodic Table with group labels and chemical families identified Actinides Note: Placement of elements 58 – 71 and 90 – 103 saves space

  23. Representative/Main Group Elements A groups—Longer columns • Alkali Metals • 1A= first group • Very reactive • All are metals except for H • Tend to form +1ions • React with oxygen • Form compounds that dissolve in water • Yield strongly caustic or alkaline solution (Na2O)

  24. Representative/Main Group Elements A groups—Longer columns • Alkaline Earth Metals • 2A= second group • Reactive • Tend to form +2ions • Oxygen compounds are strongly alkaline (MgO) • Many are not water soluble

  25. Representative/Main Group Elements A groups—Longer columns • Halogens • 7A= next to last group on right • Reactive • Form diatomic molecules in elemental state • 2 gases – F2, Cl2 • 1 liquid – Br2 • 2 solids – I2, At2 • Form–1ions with alkali metals—salts (e.g. NaF, NaCl, NaBr, and NaI)

  26. Representative/Main Group Elements A groups—Longer columns • Noble Gases • 8A = last group on right • Inert—very unreactive • Only heavier elements of group react and then very limited • Don’t form charged ions • Monatomic gases (e.g. He, Ne, Ar)

  27. Transition Elements B groups—shorter columns • All are metals • In center of table • Begin in fourth row • Tend to form ions with several different charges e.g. • Fe2+ and Fe3+ • Cu+ and Cu2+ • Mn2+, Mn3+, Mn4+, Mn5+, Mn6+, and Mn7+ Note: Last 3 columns all have 8B designation

  28. Inner Transition Elements • At bottom of periodic table • Tend to form +2 and +3 ions Lanthanide elements • Elements 58 – 71 Actinide elements • Elements 90 – 103 • All actinides are radioactive

  29. Metals, Nonmetals, or Metalloids • Elements break down into three broad categories • Organized by regions of periodic table Metals • Left-hand side • Sodium, lead, iron, gold Nonmetals • Upper right hand corner • Oxygen, nitrogen, chlorine Metalloids • Diagonal line between metals and nonmetals • Boron to astatine

  30. Metals, Nonmetals, or Metalloids

  31. Metals • Most elements in periodic table Properties • Metallicluster • Shine or reflect light • Malleable • Can be hammered or rolled into thin sheets • Ductile • Can be drawn into wire • Hardness • Some hard – iron and chromium • Some soft – sodium, lead, copper

  32. Properties of Metals • Conductheat and electricity • Solids at room temperature • Melting points (mp) > 25 °C • Hg only liquid metal (mp = –39 °C) • Tungsten (W) (mp = 3400 °C) • Highest mp for a metal • Chemicalreactivity • Varies greatly • Au, Pt very unreactive • Na, K very reactive

  33. Nonmetals • Seventeen elements • Upper right hand corner of periodic table • Exist mostly as compounds rather than as pure elements • Many are gases • Monatomic (Noble) He, Ne, Ar, Kr, Xe, Rn • Diatomic H2, O2, N2, F2, Cl2 • Some are solids: I2, Se8, S8, P4, C • Three forms of carbon (graphite, coal, diamond) • One is liquid: Br2

  34. Properties of Nonmetals • Brittle • Pulverize when struck • Insulators • Non-conductors of electricity and heat • Chemical reactivity • Some inert • Noble gases • Some reactive • F2, O2, H2 • React with metals to form ionic compounds

  35. Metalloids • EightElements • Located on diagonal line between metals and nonmetals • B, Si, Ge, As, Sb, Te, Po, At Properties • Between metals and nonmetals • Metallic shine • Brittle like nonmetal • Semiconductors • Conduct electricity • But not as well as metals • Silicon (Si) and germanium (Ge)

  36. Your Turn! Which of the following statements is correct? • Cu is a representative transition element • Na is an alkaline earth metal • Al is a metalloid in group 3A • F is a representative halogen • None of these are correct

  37. Your Turn! All of the following are characteristics of metals except: • Malleable • Ductile • Lustrous • Good conductors of heat • Acts as a semiconductor

  38. Ions and Ionic Compounds Ions • Transfer of one or more electrons from one atom to another • Form electrically charged particles Ionic compound • Compound composed of ions • Formed from metal and nonmetal • Infinite array of alternating Na+ and Cl– ions Formula unit • Smallest neutral unit of ionic compound • Smallest whole-number ratio of ions

  39. Formation of Ionic Compounds Metal + Non-metal  ionic compound 2Na(s) + Cl2(g)  2NaCl(s)

  40. Cations Positively charged ions Formed from metals Atoms lose electrons e.g. Na has 11 e– and 11 p Anions Negatively charged ions Formed from non-metals Atoms gain electrons e.g. Cl has 17 e– and 17 p Ionic Compounds Na+has 10 e– and 11 p Cl–has 16 e– and 17 p

  41. Experimental Evidence for Ions Electrical conductivity requires charge movement Ionic compounds: • Do not conduct electricity in solid state • Do conduct electricity in liquid and aqueous states where ions are free to move Molecular compounds: • Do not conduct electricity in any state • Molecules are comprised of uncharged particles

  42. Ions of Representative Elements • Can use periodic table to predict ion charges • When we use North American numbering of groups: Cation positive charge = group number

  43. Noble gases are especially stable Nonmetals Negative(–) charge on anion = number of spaces you have to move to right to get to noble gas Expected charge on O is Move two spaces to right O2– What is expected charge on N? Move three spaces to right N3 – Ions of Representative Elements

  44. Cation given first in formula Subscripts in formula must produce electrically neutral formula unit Subscripts must be smallest whole numbers possible Divide by 2 if all subscripts are even May have to repeat several times Charges on ions not included in finished formula unit of substance If no subscript, then 1 implied Rules For Writing Ionic Formulas

  45. Determining Ionic Formulas Ex. Formula of ionic compound formed when magnesium reacts with oxygen • Mg is group 2A • Forms +2 ion or Mg2+ • O is group 6A • Forms –2 ion or O2– • To get electrically neutral particle need • 1:1 ratio of Mg2+ and O2– • Formula: MgO

  46. Determining Ionic Formulas “Criss-cross” rule • Make magnitude of charge on one ion into subscript for other • When doing this, make sure that subscripts are reduced to lowest whole number. Ex. What is the formula of ionic compound formed between aluminum and oxygen ions? Al3+ O2– Al2O3

  47. Your Turn! Which of the following is the correct formula for the formula unit composed of potassium and oxygen ions? • KO • KO2 • K2O • P2O3 • K2O2

  48. Your Turn! Which of the following is the correct formula for the formula unit composed of Fe3+ and sulfide ions? • FeS • Fe3S2 • FeS3 • Fe2S3 • Fe4S6

  49. Cations of Transition Metals Transition metals • Center (shorter) region of periodic table • Much less reactive than group 1A and 2A • Still transfer electrons to nonmetals to form ionic compounds • number of electrons transferred less clear • Form more than one positive ion • Can form more than one compound with same non-metal e.g. Fe + Cl FeCl2 and FeCl3

  50. Cations of Post-transition Metals Post-transition metals • Nine metals Ga, In, Sn, Tl, Pb, Bi, Uut, Uuq, Uub • After transition metals and before metalloids • Two very important ones – tin (Sn) and lead (Pb) • Both have two possible oxidation states • Both form two compounds with same nonmetal e.g. Ionic compounds of tin and oxygen are • SnO and SnO2 • Bismuth • Only has +3 charge • Bi3+

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