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Chapter 2. Atoms Molecules & Ions KEY CONCEPTS

Chapter 2. Atoms Molecules & Ions KEY CONCEPTS. Chemical laws Dalton's atomic theory Atom Molecules Ions Average atomic mass Periodic table Formula of a substance Name of a Substance

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Chapter 2. Atoms Molecules & Ions KEY CONCEPTS

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  1. Chapter 2. Atoms Molecules & Ions KEY CONCEPTS • Chemical laws • Dalton's atomic theory • Atom • Molecules • Ions • Average atomic mass • Periodic table • Formula of a substance • Name of a Substance • Molecular or Covalent compounds • Ionic compounds

  2. Alchemist • Early scientist observed changes of matter. They called these changes chemical reactions when there are changes in substances or the physical properties of the matter. They also observed a pattern or a repeatable observation in chemical reactions.

  3. Three Chemical Laws: • Law of Conservation of Mass: • Law of Constant Proportions: • Law of Multiple Proportions:

  4. Law of Conservation of Mass • Total mass after a chemical reaction is same as before the reaction. • H2 + 1/2 O2 ---> H2O • Hydrogen (4g) + oxygen (32g) ----> water 36g after the reaction.

  5. Law of Constant (Definite) Proportions A given chemical compound always contains exactly the same proportion of elements by weight. 36gofwater contains 4g of hydrogen and 32g of oxygen take any other chemical compund.

  6. Law of Multiple Proportions When two elements make a series of chemical compounds, the ratio of the masses of the second elements that combine with 1 gram of the first element can always be reduced to simple whole numbers. C O E.g. carbon monoxide 1g 1.33g carbon dioxide 1g 2.66g

  7. Dalton’s atomic theory • All matter is composed of atoms -- the smallest particle of an element that takes part in a chemical reaction. • All atoms of an element are alike. • Compounds are combinations of atoms of one or more elements. The relative number of atoms each element is always the same. • Atoms cannot be created or destroyed by a chemical reaction. They only change how they combine with each other.

  8. Models of matter • Models are commonly used to help visualize atoms and molecules. • Atom - The smallest unit of an element that has all of the properties of an element. • Molecule - The smallest unit of a pure substance that has the properties of that substance. • It may contain more that one atom and more than one element. • Ions - Charged particles formed by the transfer of electrons between atoms or molecules

  9. What is an Atom? • Very small particle. • Smallest particles of elements and molecules • There about 110 types of elements or Atoms. • Different atoms have different physical properties and chemical reactivity

  10. Structure of the atom • Atoms have a specific arrangement. • Nucleus Small, dense, positively charged • in the center of an atom. • Electrons Surround the nucleus. • Negative charge.

  11. What are these? a) atoms b) nucleus c) electrons d) protons e) neutrons f) isotope g) atomic number(Z)

  12. Atomic Symbols • Each element is assigned a unique symbol. • arsenic As potassium K • barium Ba nickel Ni • carbon C nitrogen N • chlorine Cl oxygen O • hydrogen H radon Rn • helium He titanium Ti • gold Au uranium U • Each symbol consists of 1 or 2 letters. The first is capitalized and the second is lower case. • Symbol may not match the name - often had a different name to start with. • Some elements (about 11) the names were not in English. E.g., Sodium-Na (natrium-latin), • potassium-K(kalium-latin).

  13. Atomic masses • Atoms are composed of protons, neutrons and electrons. • Almost all of the mass of an atom comes from the protons and neutrons. • All atoms of the same element will have the same number of protons. The number of neutrons may vary - isotopes. • Most elements exist as a mixture of isotopes.

  14. Isotopes • *Atoms of the same element but having different masses. • *All isotopes of an element have same atomic number • *Each isotope has a different number of neutrons. • Isotopes of hydrogen H H H • Isotopes of carbon C C C 1 1 2 1 3 1 13 6 14 6 12 6

  15. 31 15 138 56 238 92 P Ba U The atomic symbol & isotopes • Isotopic symbol: atomic symbol showing atomic number (Z) and mass number (A) • Determine the number of protons, neutrons and electrons in each of the following. 2+

  16. Isotopes • Most elements occur in nature as a mixture of isotopes. • Element Number of stable isotopes • H 2 • C 2 • O 3 • Fe 4 • Sn 10 • This is one reason why atomic masses are not whole numbers. They are based on averages.

  17. Average atomic masses • Most elements exits as a mixture of isotopes. • Each isotope may be present in different amounts. • The masses listed in the periodic table reflect the world-wide average for each isotope. • One can calculate the average atomic weight of an element if the abundance of each isotope for that element is known.

  18. Average atomic masses • Example. • Silicon exists as a mixture of three isotopes. Determine it’s average atomic mass based on the following data. • Isotope Mass (u) Abundance • 28Si 27.976 9265 92.23 % • 29Si 28.976 4947 4.67 % • 30Si 29.973 7702 3.10 %

  19. What is a mass spectrometer? • Mass spectrometer measures masses of different isotopes of an elements and their fractional or percentages abundance. (figure 7.7, page 283) • As a reference point, we use the • atomic mass unit (u) - 1/12th mass of a 12C atom.

  20. Diagram of mass (positive ion) spectrometer

  21. How do you calculate average Atomic Mass? • Ma x a + Mb x b • ------------------------ = AAM • 100 • Ma = mass of isotope a • Mb = mass of isotope b • a = percent abundance of a • b = percent abundance of b • AAM = Average atomic mass (Reported • on the Periodic Table)

  22. How do you calculate average Atomic Mass? Ma x a + Mb x b = AAM Ma = mass of isotope a Mb = mass of isotope b a = fractional abundance of a b = fractional abundance of b AAM = Average atomic mass (Reported on the Periodic Table)

  23. Gallium in nature consists of two isotopes, gallium-69, with a mass of 68.926 u and a fractional abundance of 0.601; and gallium-71, with a mass of 70.925 u and a fractional abundance of 0.399. • Calculate the weighted average atomic mass of gallium. • 1) Ma x a + Mb x b = AAM • Ma x a(%) + Mb x b(%) • 2) ----------------------------------- = AAM • 100

  24. Ma (69Ga ) =68.926 u, • a = percent abundance of 69Ga = 0.601 x 100 • Mb (71Ga ) = 70.925 u, • b = percent abundance of 71Ga = 0.399 x 100 • We can obtain an equation with one unknown, AAM. • AAM = 68.926x(0.601 x 100)+70.925 x(0.399x100) • 100 • AAM (Ga) = 4142.5 + 2829.9 • 100 • AAM (Ga) = 6972.3 = 69.723 • 100 • AAM (Ga) = 69.7 u (amu)

  25. Periodic Table • Periodic table is an arrangement of all known element according to their atomic number and chemical properties.

  26. Who is Dmitri Mendeleev? • Mendeleev, Dmitri (1834-1907): Russian chemist • Mendeleev is best known for • his work on the periodic table; • arranging the 63 known • elements into a Periodic • Table based on Atomic Mass

  27. Dimitri Mendeleev created this, the original, periodic table.

  28. La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Modern periodic table 1 2 13 14 15 16 17 18 I A II A III A IV A V A VI A VIIA 0 H He 1 2 3 4 5 6 7 Li Be B C N O F Ne 3 4 5 6 7 8 9 10 11 12 III B IVB V B VIB VIIB VIII B IB IIB Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe * Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn + Fr Ra Lr * +

  29. 47 Silver Ag 107.87 Information that may be in the table Atomic number Name of the element Elemental Symbol Atomic mass

  30. Vertical columns- groups,families • Horizontal columns- periods • Elements in a group have similar • chemical properties group IA - alkali metal: Li, Na, K Rb, Cs, Fr group IIA- alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra group VIIA - Halogens: Cl, Br, I, At group 0 - Noble gases: He, Ne, Ar, Kr, Xe, Rn

  31. A group or family Groups are assigned Roman numerals with an A or B I A II A III A IV A V A VI A VIIA 0 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar III B IVB V B VIB VIIB VIII IB IIB K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

  32. A row or period Periods are assigned numbers H He 1 2 3 4 5 6 7 Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

  33. Common group names I A II A III A IV A V A VI A VIIA 0 He H Li Be B C N O F Ne Na Mg Al Si P S Cl Ar III B IVB V B VIB VIIB VIII B IB IIB K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

  34. Solid Liquid Gas Elemental states atroom temperature H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe * Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn + Fr Ra Lr * La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No +

  35. The known elements • 112 elements are currently known • 89 are metals • 31 are radioactive • 22 are synthetic (all radioactive) • 11 occur as gases • 2 occur as liquids • Let’s take a look at them on the table.

  36. What are these? • Transition Metals • Actinides • Lanthanides • Semimetals or Metalloids • Ionic Charges • Poly atomic ions and their charges

  37. Naming Chemical Compounds • Binary compunds: Combinations of two differerent elements • Common names: Water-H2O • Vinger, Ammomia, Chloroform • Systematic Names: Names given starting from • element names and/or following certain rules • Naming depend on the type of compound • Ionic Compounds: • Name:Sodium chloride • Formula:NaCl • Molecular or Covalent Compounds: • Carbon tetrachloride- CCl4

  38. Types of Compounds • Ionic compounds: • Metal + non-metal • Type I : ionic compound • (fixed charge) NaCl • Type II ionic compound • FeCl2 and FeCl3, SnCl2 and SnCl4 • Covalent Compounds: • non-metal + non-metal • sulfur dioxide: SO2

  39. Formula • Formula are used to represent elements and compound. • For molecular compounds, formula tell how many of each kind of atom are in a molecule. • For ionic compounds, formula tell the simples ratio of cations and anions. Molecular Weight ? and Formula Weight?

  40. What is Empirical Formula • Simple whole number ratio of each atom expressed in the subscript of the formula. • Molecular Formula = C6H12O6 of glucose • Empirical Formula = CH2O • Emiprical formula is calculated from % composition

  41. Ions • Ions are charged particles formed by the transfer of electrons between elements or combinations of elements. • Cation - a positively charged ion. • Ca Ca2+ + 2e- • Anion - a negatively charged ion. • F2 + 2e- 2F-

  42. Cations Al3+ Na+ Mg2+ O2- N3- Cl- Anions Ionic compounds Some simple ions Formula for some ionic compounds NaCl MgCl2 AlCl3 Na2O MgO Al2O3 Na3N Mg3N2 AlN

  43. Naming Ionic compounds • When an element forms only one compound with a given anion. • name the cation • name the anion using the ending (-ide) • NaCl sodium chloride • MgBr2 magnesium bromide • Al2O3 aluminum oxide • K3N potassium nitride

  44. Metals with multiple charges • Transition metals. • Here it is easier to list the ones that have a single common oxidation state. • All Group 3B - 3+ • Ni, Zn, Cd - 2+ • Ag - 1+ • Lanthanides and actinides - 3+

  45. Ionic compounds: • a) Type I ionic compound • NaCl, CaCl2, AlCl3 • sodium chloride, calcium chloride, • Aluminum chloride • b) Type II ionic compound • FeCl2 and FeCl3 • SnCl2 and SnCl4 • iron(II)chloride, iron(III)chloride, tin(II)chloride, tin(IV)chloride

  46. Examples FeCl2 iron(II) chloride FeCl3 iron(III) chloride SnS tin(II) sulfide SnS2 tin(IV) sulfide AgCl silver chloride CdS cadmium sulfide • Note • Some transition metals only have a single state • so the roman numeral may be omitted. • Some main group metals, with high atomic number • have more than one state, roman numbers are used.

  47. Naming Hydrates • Hydrates are substances that include water into their formula. • The water is not actually part of the chemical substance and this is reflected in the way the formula is written. • Example: CuSO4 . 5 H2O

  48. Example • Name the following compounds • Na2S • \______ sodium sulfide • HCl • \______ hydrogen chloride • CaH2 • \______ calcium hidride • SrO • \______ strontium oxide

  49. Polyatomic ions • Name the following • AgNO3 • \______ silver nitrate • Mg(C2H3O2)2 • \_______ magnesium acetate • Fe2(SO4)3 • \____ iron(III) sulfate

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