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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 & 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 & early 1900s • Cathode ray tube experiments showed that atoms are made up of subatomic particles • Discovered negatively charged particles moving from • Cathode – negative electrode to • 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 x 1019 C • Combined with Thomson’s experiment to get mass of electron • m = 9.09 x 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 1800x mass of electron • 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+) & neutrons(1n) are located

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

  9. Discovery of Neutron (1n) • First postulated by Rutherford & 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 • Caused free neutron to be created

  10. Properties of Subatomic Particles • 3 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= # 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 = (# protons) + (#neutrons) • A = Z + N • For charge neutrality, number of electrons & protons must be equal Atomic Symbols • Summarize information about subatomic particles • Every isotope defined by 2 numbers Z & A • Symbolized by Ex.What is the atomic symbol for helium? He has 2 e–, 2 n & 2 p+ Z = 2, A = 4

  13. Isotopes • Most elements are mixtures of 2 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): Ex. • 3 isotopes of hydrogen (H) • 4 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. Learning Check: • Fill in the blanks: symbol neutrons protons electrons 60Co 81Br 36 29 29 33 27 27 46 35 35

  16. 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

  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 Ex. 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 & 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 * 23.9850 u = 18.946 u 24Mg 0.1000 * 24.9858 u = 2.4986 u 25Mg 0.1101 * 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) & 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 & 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 Metals except for H • Tend to form +1ions • React with oxygen • Form compounds that dissolve in water • Yield strongly caustic or alkaline solution (M2O)

  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 (MO) • Many are not water soluble • Accumulate in ground

  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 • 1 liquid • 2 solids • Form–1ions with alkali metals—salts

  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 & then very limited • Don’t form charged ions • Monatomic gases

  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 Ex. • Fe2+ and Fe3+ • Cu+ and Cu2+ • Mn2+, Mn3+, Mn4+, Mn5+, Mn6+, Mn7+ Note: Last 3 columns all have 8B designation

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

  29. Metals, Nonmetals, or Metalloids • Elements break down into 3 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 & 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 & chromium • Some soft – sodium, lead, copper

  32. Properties of Metals • Conduct heat & electricity • Solidsat Room Temperature • Melting points (mp) > 25 °C • Hg only liquid metal (mp = –39 °C) • Tungsten (W) (mp = 3400 °C) • Highest known for metal • Chemicalreactivity • Varies greatly • Au, Pt very unreactive • Na, K very reactive

  33. Nonmetals • 17 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 • 3 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 • 8Elements • Located on diagonal line between metals & nonmetals • B, Si, Ge, As, Sb, Te, Po, At Properties • Between metals & nonmetals • Metallic shine • Brittle like nonmetal • Semiconductors • Conduct electricity • But not as well as metals • Silicon (Si) & 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 semimetal in group IIIA • 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 • Tend to gain electrons in chemical reactions

  38. Ions & Ionic Compounds Ions • Transfer of 1 or more electrons from 1 atom to another • Form electrically charged particles Ionic compound • Compound composed of ions • Formed from metal & nonmetal • Infinite array of alternating Na+ & 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 Ex. Na has 11 e– & 11 p+ Anions Negatively charged ions Formed from non-metals Atoms gain electrons Ex. Cl has 17 e– & 17 p+ Ionic Compounds Na+has 10 e– & 11 p+ Cl–has 16 e– & 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 & 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 #

  43. Noble gases are especially stable Nonmetals Negative() charge on anion = # spaces you have to move to right to get to noble gas Expected charge on O is Move 2 spaces to right O2– What is expected charge on N? Move 3 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+ & 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 & 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 & 2A • Still transfer electrons to nonmetals to form ionic compounds • # of electrons transferred less clear • Form more than 1 positive ion • Can form more than 1 compound with same non-metal Ex. Fe + Cl FeCl2 & FeCl3

  50. Cations of Post-transition Metals Post-transition metals • 9 metals Ga, In, Sn, Tl, Pb, Bi, Uut, Uuq, Uub • After transition metals & before metalloids • 2 very important ones – tin (Sn) & lead (Pb) • Both have 2 possible oxidation states • Both form 2 compounds with same nonmetal Ex. Ionic compounds of tin & oxygen are • SnO & SnO2 • Bismuth • Only has +3 charge • Bi3+

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