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Chapter 2

Chapter 2 . Atoms , Molecules, and Ions. History of the Atom. Democritus (400 BC). Proposed that matter was composed of tiny, invisible particles. Gr. atomos. Aristotle. Said that all substances are composed of four elements: Earth, Wind, Fire, and Water. John Dalton (1803 AD).

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Chapter 2

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  1. Chapter 2 Atoms, Molecules, and Ions

  2. History of the Atom

  3. Democritus (400 BC) • Proposed that matter was composed of tiny, invisible particles. • Gr. atomos

  4. Aristotle • Said that all substances are composed of four elements: Earth, Wind, Fire, and Water.

  5. John Dalton (1803 AD) • the first coherent atomic theory • Matter is composed of small particles called atoms. • All atoms of an element are identical. • During chemical reactions, atoms are neither created nor destroyed, but are simply rearranged.

  6. J.J. Thompson (1897 AD) • Used cathode rays to calculate the atomic mass of atoms.

  7. Cathode Ray Tube Experiment

  8. Ernest Rutherford (1910 AD) • alpha particles (He nuclei) were shot at a thin piece of gold foil. • most of the particles passed through the foil with little or no deflection, some were deflected to a great degree.

  9. Gold Foil Experiment

  10. History 400 -370 BC - Democritus thought that there must be atoms, “invisible particles”. 384-322 BC – Aristotle refused this theory. 1700 – Isaac Newton again favored the idea of smaller invisible particles. 1800 – John Dalton formed the atomic theory.

  11. Dalton’s Atomic Theory • Elements composed of small particles called atoms. • All atoms of the same element are identical in physical properties but different from atoms of other elements. • Atoms of one element can not change into atoms of different elements with chemical reactions. • Compounds are composed of atoms of different elements and are consistent in number and type of elements.

  12. Dalton’s Atomic Theory Vocabulary • Atom – the smallest particle of an element that retains the chemical identity of the element. • Compound – contains atoms of two or more elements. • Conservation of matter – atoms can neither be created nor destroyed in a chemical reaction.

  13. Continuing History 1897 – J. J. Thomson – using a cathode ray determined the presence of negative particles, electrons, and the “plum pudding” model. 1911 –Ernest Rutherford – using alpha particles through gold foil determined electrons were not evenly spaced and determined the presence of a nucleus. 1919 – Rutherford- determined the presence of protons. 1932 – James Chadwick – determined the presence of neutrons.

  14. Modern Atomic Structure Atoms consist of subatomic particles:

  15. Atomic Number The number of protons in an atom of an element. • Each element has a different atomic number or number of protons. • Each element has no charge. • Each element has the same number of electrons as protons to keep neutral.

  16. Periodic Table of Elements • 1750 only 17 elements • 1800 – 31 elements • 1865 – 63 elements • Today – 117 elements • Antoine Lavoisier – categorized elements into metals, nonmetals, gases, and earths • Dmitri Mendeleev (Russia)- 1865 – categorized 63 elements according to atomic weight along with Lothar Meyer (Germany).

  17. Mendeleev • Modeled chart of elements after the solitaire card game. • Arranged the elements into rows in the order of increasing mass so that elements with similar properties were in the same column like suits in the card game. • Within columns, atomic masses increased from top to bottom leading to the periodic table.

  18. Mendeleev • Since many elements were still undiscovered, he left gaps in the chart where he thought the undiscovered elements should be. • The structure of the table lead to the prediction and discovery of gallium which had similar properties as aluminum. • 1913 Rutherford’s nuclear model of the atom lead to atomic numbers, verification of Mendeleev’s table.

  19. Periodic Table • Organized by increasing atomic number. • Basic Info: 17 = atomic number Cl = symbol • 35.453 = atomic mass 17 Cl 35.453

  20. Organization Table is configured into: • Periods – rows on the periodic table. • Groups – columns on the table with elements in the same group having similar physical properties. Further organized into: • Metals • Metalloids • Nonmetals

  21. Organization Metals • good conductors of heat and electricity • Malleable (hammered into thin sheets) • Ductile (drawn into thin wires) • Lustrous (shiny) Nonmetals • Poor conductors • Mostly gases • If metal then brittle Metalloid • Demonstrate both metal and nonmetal properties

  22. Organization Groups • Group 1 – Alkali metals • Group 2 – Alkaline earth metals • Group 7 – Halogens • Group 8 – Noble gases • Group B – Transition metals

  23. Organization Alkali metals – very reactive with water and oxygen. They have low densities and melting points. They all have 1 valence electron so readily give away 1 electron in s orbital. Ex. Sodium and potassium react violently with water such that they will react with the water in human skin igniting the hydrogen molecules and burn the skin.

  24. Organization Alkaline earth metals – have 2 valence electrons. Differences in reactivity along these elements is shown by the ways they react with water. More dense and higher melting temperatures Ex. Calcium, strontium and barium react easily with water. Magnesium reacts with hot water. Beryllium has no reaction in water.

  25. Organization Halogens – highly reactive with metals. They all have 7 valence electrons. Nobel gases – mostly nonreactive colorless, odorless gas that give off different colors when excited. Ex. Helium-pink Neon-orange/red Argon-lavender Krypton-white Xenon-blue

  26. Atomic & Ionic Radii Trends • Columns of periodic table – radius increases from top to bottom. • Periods of periodic table – radius decreases from left to right. • Cations have smaller radii than parent elements. • Anions have larger radii than parent elements.

  27. Molecule Assembly of two or more atoms tightly bound together with no net charge. Chemical formula - representation of the number and type of elements in a compound. H2O, CO2, CH4 Molecular formula– actual chemical formula of a molecule indicating the actual number of molecules of each atom Empirical formula – simplified formula of a molecule indicating the smallest ratio of atoms of elements.

  28. Writing Chemical Formulas • Each atom (element) is represented by it symbol. • The number of atoms of each element is represented by a subscript. • When the number of atoms is 1 then the subscript is not written and understood to be one.

  29. Chemical Formulas Examples of formulas • Ca3(PO4)3 • Al(NO3)3 • H2SO4 • 3CO2

  30. Diatomic molecules Molecules that exist is pairs • Hydrogen • Nitrogen • Oxygen • Fluorine • Chlorine • Bromine • Iodine

  31. Ions A charged entity produced by taking a neutral atom and adding or removing one or more electrons. • Cation-positively charged particle due to the removal of an electron. K+ , Mg2+ • Anion-negatively charged particle due to the acceptance of an electron (tend to be nonmetals). Cl- , S2-

  32. Ions How many electrons are contained in each of the following ions? • Ba2+ • P3- • Sn2+ • Cl-

  33. Writing Formulas • There must be both positive ions (cations) and negative ions (anions) present. • The numbers of cations and anions must have a net charge of zero or the sum of the oxidation numbers is zero. • Cation is first then the anion.

  34. Writing Formulas Ex. Na and Cl Na+ Cl- → NaCl Charge: +1 Charge: -1 Net = 0 Oxidation number = charge Ex. Mg and Cl Mg2+ Cl - Cl- →MgCl2 Charge: +2 2 x (-1) net =0

  35. Writing Formulas • Calcium and chlorine • Sodium and sulfur • Lithium and nitrogen • Phosphorus and Calcium • Barium and oxygen • Sulfur and Aluminum • Potassium and phosphorus

  36. Writing Formulas • Ionic Compound – a molecule that contains both a metal and a nonmetallic elements. NaCl MgCl2 • Molecular Compound – a molecule that contains only nonmetals. CO2 CH4 • Polyatomic ions – atoms joined as a molecule with a charge.

  37. Writing Formulas • Iron (II) and chlorine • Cooper (I) and fluorine • Calcium and hydroxide • Chromium (II) and peroxide • Hydrogen and oxygen • Hydrogen and phosphate

  38. Nomenclature Naming cations: • Metals with single charge – cation is metal name with “ion” added • Metal with multiple charge – cation is metal name with charge indicated by Roman numeral and “ion” added. • Polyatomic – nonmetal name with “ium” replacing end.

  39. Nomenclature Naming anions: • Monatomic – element name with end replaced with “ide” . • Polyatomic – • With oxygen (oxyanion)– element name with end replaced with “ate”. • With hydrogen & oxyanion – polyatomic name prefix “hydrogen” added.

  40. Nomenclature • Naming ionic compounds: Cation is named first using cation name then anion name. • Naming molecular compounds: • Name element further to left first. • If same group, name higher atomic number first. • Anion is named using anion name (ide) • Atoms are indicated by following prefixes: • One – mono six - hexa • Two – di seven - hepta • Three – tri eight - octa • Four – tetera nine - nona • Five - penta ten - deca

  41. Nomenclature Naming acids: • If anion name ends in “ide” change to “ic” and add “hydrogen” to name. These are molecules with no oxygen. • If anion name ends in “ate” change to “ic” and add “acid”. • If anion name ends in “ite” change to “ous” and add “acid”.

  42. Is it Ionic or molecular? Naming Flow Chart Ionic Molecular nonmetals Transition metal With multiple charges? Use prefixes to tell number of each atom present Yes Metal – roman numeral Nonmetal – “ide” ending Or polyatomic ion name No Metal – metal name Nonmetal – “ide” ending Or polyatomic ion name

  43. Isotopes

  44. Radioactive Decay

  45. C14 Decay

  46. Carbon Dating

  47. Carbon Dating Video Click on the picture above to view.

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