1 / 56

Atoms, Ions, and Molecules: Matter Starts Here

2. Atoms, Ions, and Molecules: Matter Starts Here. Chapter Outline. 2.1 The Nuclear Model of Atomic Structure Electrons Radioactivity and the Nuclear Atom Protons and Neutrons 2.2 Isotopes 2.3 Average Atomic Mass 2.4 The Periodic Table of the Elements 2.5 Trends in Compound Formation

robert-harrell
Download Presentation

Atoms, Ions, and Molecules: Matter Starts Here

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. 2 Atoms, Ions, and Molecules: Matter Starts Here

  2. Chapter Outline • 2.1 The Nuclear Model of Atomic Structure • Electrons • Radioactivity and the Nuclear Atom • Protons and Neutrons • 2.2 Isotopes • 2.3 Average Atomic Mass • 2.4 The Periodic Table of the Elements • 2.5 Trends in Compound Formation • 2.6 Naming Compounds and Writing Formulas • 2.7 Nucleosynthesis

  3. Atomic Structure: Electrons • J. J. Thomson (1897) • Beam from cathode ray tube is deflected toward positively charged plate. • Atoms contain negatively charged particles with a constant mass-to-charge ratio.

  4. Mass of an Electron • Robert Millikan (1909) • Determined the mass and charge of an electron with his oil-drop experiment. • e– = –1.602 x 10–19 C • me = 9.109 x 10 –28 g

  5. Plum-pudding model: e– distributed throughout diffuse, positively charged sphere Thomson’s Model of the Atom

  6. Radioactivity and the Nuclear Atom • Henri Becquerel (1896) • Some materials produce invisible radiation, consisting of charged particles. • Radioactivity • Spontaneous emission of high energy radiation and particles • Beta particles (, high energy electrons) • Alpha particles (, +2 charge, mass = He nucleus)

  7. Atomic Structure: The Nucleus • Rutherford’s experiment: • Bombard a thin gold foil with  particles to test Thomson’s model of the atom.

  8. Rutherford’s Experiment (c) Actual results (b) Expected results from plum-pudding model

  9. The Nuclear Atom • The nucleus: • is the positively charged center of an atom, containing nearly all of the atom’s mass. • is about 1/10,000 the size of the atom. • consists of two types of particles: • Protons: positively charged subatomic particles • Neutrons: electrically neutral subatomic particles

  10. The Nuclear Atom (cont.) Nucleus has protons (+ charge) plus neutrons (neutral).

  11. Atomic Mass Units • Atomic mass units (amu) • Unit used to express the relative masses of atoms and subatomic particles • Equal to 1/12 of a carbon atom: • 6 protons + 6 neutrons = 12 amu • 1 amu = 1 Dalton (Da)

  12. The Nuclear Atom: Summary

  13. Chapter Outline • 2.1 The Nuclear Model of Atomic Structure • 2.2 Isotopes • Positive Ray Analyzer Experiments • Isotopes: Atoms of the Same Element, but Different Masses • 2.3 Average Atomic Mass • 2.4 The Periodic Table of the Elements • 2.5 Trends in Compound Formation • 2.6 Naming Compounds and Writing Formulas • 2.7 Nucleosynthesis

  14. Aston’s Positive-Ray Analyzer Ne gas ions of different masses strike the detector in different locations.

  15. Aston’s Experiment • Positive-ray analyzer results: • Two different kinds of neon gas atoms exist. • 90% = 20 amu • 10% = 22 amu • Aston proposed theory of isotopes. • Isotopes • Atoms of an element containing the same # of protons but different numbers of neutrons.

  16. Symbols of Isotopes A X Z Atomic Mass (A) – total number of nucleons (protons + neutrons) in the nucleus. Elemental Symbol – a one- or two-letter symbol to identify the type of atom. Atomic Number (Z) – the number of protons in the nucleus; determines the identity of the element.

  17. Practice: Isotopic Symbols • Use the format AX to write the symbol for the nuclides having 26 protons and 30 neutrons. - Collect and Organize: We have a nuclide with 26 protons and 30 neutrons.

  18. Practice: Isotopic Symbols • Use the format AX to write the symbol for the nuclides having 26 protons and 30 neutrons. - Analyze: The number of protons represents the atomic number and identity of the element, and the total number of nucleons (protons + neutrons) represents the atomic mass. 2-18

  19. Practice: Isotopic Symbols • Use the format AX to write the symbol for the nuclides having 26 protons and 30 neutrons. - Solve: The number of protons (26) identifies this nuclide as Fe (iron). The total number of nucleons is 26 + 30 = 56 and is the mass of the nuclide. So, the symbol for this nuclide would be 56Fe. 2-19

  20. Practice: Isotopic Symbols • Use the format AX to write the symbol for the nuclides having 26 protons and 30 neutrons. - Think About It: The mass of Fe in the periodic table is 55.85, which is close to the mass of this nuclide, so this represents a reasonable result for an isotope of iron. 2-20

  21. Chapter Outline • 2.1 The Nuclear Model of Atomic Structure • 2.2 Isotopes • 2.3 Average Atomic Mass • Weighted Average of Isotopic Masses • Natural Isotopic Abundances • 2.4 The Periodic Table of the Elements • 2.5 Trends in Compound Formation • 2.6 Naming Compounds and Writing Formulas • 2.7 Nucleosynthesis

  22. Average Atomic Masses • Average atomic mass: • Weighted average of masses of all isotopes of an element • Calculated by multiplying the natural abundance of each isotope by its mass in amu and then summing these products • Natural abundance: • proportion of a particular isotope; usually expressed as a percentage relative to all the isotopes for that element in a natural sample

  23. Weighted Average Example Neon has three naturally occurring isotopes. • Average atomic mass of neon: (19.9924 x 0.904838) + (20.99395 x 0.002696) + (21.9914 x 0.092465) = 20.1797 amu

  24. Chapter Outline • 2.1 The Nuclear Model of Atomic Structure • 2.2 Isotopes • 2.3 Average Atomic Mass • 2.4 The Periodic Table of the Elements • Periodicity and Mendeleev’s Table of the Elements • Navigating the Modern Periodic Table • 2.5 Trends in Compound Formation • 2.6 Naming Compounds and Writing Formulas • 2.7 Nucleosynthesis

  25. Dmitri Mendeleev (1872) Ordered elements by atomic mass Arranged elements in columns based on similar chemical and physical properties Left open spaces in the table for elements not yet discovered Mendeleev’s Periodic Table

  26. The Modern Periodic Table Also based on a classification of elements in terms of their physical and chemical properties Horizontal rows – called periods (1 → 7) Columns – contain elements of the same family or group (1 →18) Several groups have names as well as numbers

  27. Groups of Elements

  28. Nonmetals Metalloids Metals

  29. Broad Categories of Elements • Metals (left side and bottom of the table) • Shiny solids; conduct heat and electricity; are malleable and ductile • Nonmetals (right side and top of the table) • Solids, liquids and gases; nonconductors; solids are brittle. • Metalloids (between metals/nonmetals) • Shiny solids (like metals); brittle (like nonmetals); semiconductors

  30. Groups of Elements (cont.)

  31. Chapter Outline • 2.1 The Nuclear Model of Atomic Structure • 2.2 Isotopes • 2.3 Average Atomic Mass • 2.4 The Periodic Table of the Elements • 2.5 Trends in Compound Formation • Law of Multiple Proportions • Molecular Compounds • Ionic Compounds • 2.6 Naming Compounds and Writing Formulas • 2.7 Nucleosynthesis

  32. The Composition of Compounds • Law of Multiple Proportions • The ratio of the two masses of one element that react with a given mass of another element to form two different compounds is the ratio of two small whole numbers. • Examples: • SO2, SO3 • NO, NO2 • NO2 = 22.8 g O/10 g N • NO = 11.4 g O/10 g N

  33. Molecular Compounds • Molecular compounds • Composed of atoms held together in molecules by covalent bonds • Covalent bonds • Bond between two atoms created by sharing one or more pairs of electrons • Molecular compounds are composed of nonmetals.

  34. Molecular Compounds (cont.) • Molecular formula: • Shows the number and type of atoms present in one molecule of a compound • Empirical formula: • Shows the smallest whole-number ratio of elements in a compound • Example: Glucose • Molecular formula = C6H12O6 • Empirical Formula = CH2O

  35. Ionic Compounds • Cations = ions with positive charge • Anions = ions with negative charge • Ionic compounds • Consist of charged particles (ions) formed by transfer of electrons between atoms • Ions held together by electrostatic forces

  36. Ionic Compounds (cont.) • Ionic compounds are made of a metal and a nonmetal. • Metals form cations; nonmetals form anions. • Charges on ions depend on location in the periodic table. • e.g.,Group 1 Metals = +1; Halogens = –1

  37. Ionic Compounds (cont.) • e.g., NaCl • Formula unit • Smallest electrically neutral unit of an ionic compound

  38. Charges on Ions Figure 2.17

  39. Chapter Outline • 2.1 The Nuclear Model of Atomic Structure • 2.2 Isotopes • 2.3 Average Atomic Mass • 2.4 The Periodic Table of the Elements • 2.5 Trends in Compound Formation • 2.6 Naming Compounds and Writing Formulas • Binary Molecular Compounds • Binary Ionic Compounds • Binary Compounds of Transition Metals • Polyatomic Ions • Acids • 2.7 Nucleosynthesis

  40. Naming Compounds • Binary molecular compounds (e.g., SO3) • Compounds consisting of two nonmetals • First element in the formula is named first. • S = sulfur • Second element name is changed by adding suffix -ide. • O = oxygen → oxide • Add prefixes to identify quantity of atoms (see Table 2.2). • SO3 = sulfur trioxide

  41. Rules for Using Prefixes • Do not use the prefix mono-when naming first element: • SO3 monosulfur trioxide • Prefixes ending with o- or a- are modified when used with elements beginning with vowels: • P4O10 tetraphosphorus decaoxide

  42. Binary Ionic Compounds • Binary ionic compounds consist of cations (usually metals) and anions (usually nonmetals), e.g.,MgCl2. • Cation named first using name of element. • Mg = magnesium • Anion named by adding the -ide suffix to the name of the element. • Cl = chlorine → chloride • Formulas for ionic compounds must always be neutral: Mg2+ + (Cl–) x 2

  43. Binary Ionic Compounds (cont.) FeCl2: Iron(II) chloride FeCl3: Iron(III) chloride • For metals that form cations with different charges, a Roman numeral is added to indicate the charge of the cation.

  44. Practice: Ionic Compounds Write the name of the following compounds. a) NaCl b) CrCl3 Write the chemical formula of the following compounds. a) Zinc nitride b) Copper(I) oxide

  45. Common Polyatomic Ions • Polyatomic ions • Charged group of two or more atoms joined together by covalent bonds • Oxoanions • Polyatomic anions containing oxygen in combination with one or more other elements • Examples: acetate (C2H3O2–), nitrate (NO3–), carbonate (CO32–), perchlorate (ClO4–), sulfate (SO42–) from Table 2.3, page 60

  46. Practice: Polyatomic Ions • Write the names of the following compounds. a) Cr(ClO4)3 b) NH4NO3 • Write the chemical formulas for the following compounds. a) Lithium bicarbonate b) Calcium hypobromite

  47. Naming Binary Acids • Binary acids • Contain hydrogen and a monoatomic anion (e.g., Cl–, S2–) • Most common binary acids contain halogens. (e.g., HCl, HBr) • Acid names: • The prefix “hydro-” + the halogen base name + the suffix “-ic” + the word acid. • Example: HBr is hydrobromicacid.

  48. Oxoanions & Related Acids • the corresponding acid ends in: • -ic • -ite If oxoanion name ends in: -ate -ite

  49. Practice: Naming Compounds and Acids • K2Cr2O7 • Na3N • NO2 • H2CrO4 • Sodium carbonate • Sulfurous acid • Iron(II) phosphate Identify each of the following as a molecular compound, an ionic compound, or an acid. Name or give formulas for the compounds.

More Related