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Atoms, Molecules, and Ions

Atoms, Molecules, and Ions. Chemistry Timeline #1. B.C. 400 B.C. Demokritos and Leucippos use the term "atomos”.  2000 years of Alchemy . 1500's Georg Bauer: systematic metallurgy Paracelsus: medicinal application of minerals. 1600's

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Atoms, Molecules, and Ions

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  1. Atoms, Molecules, and Ions

  2. Chemistry Timeline #1 B.C. 400 B.C. Demokritos and Leucipposuse the term "atomos” 2000 years of Alchemy • 1500's • Georg Bauer: systematic metallurgy • Paracelsus: medicinal application of minerals 1600's Robert Boyle:The Skeptical Chemist.Quantitative experimentation, identification of elements • 1700s' • Georg Stahl: Phlogiston Theory • Joseph Priestly: Discovery of oxygen • Antoine Lavoisier: The role of oxygen in combustion, law of conservation of • mass, first modern chemistry textbook

  3. Chemistry Timeline #2 • 1800's • Joseph Proust: The law of definite proportion (composition) • John Dalton: The Atomic Theory, The law of multiple proportions • Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules • Amadeo Avogadro: Molar volumes of gases • Jons Jakob Berzelius: Relative atomic masses,modern symbols for the elements • Dmitri Mendeleyev: The periodic table • J.J. Thomson: discovery of the electron • Henri Becquerel: Discovery of radioactivity • 1900's • Robert Millikan: Charge and mass of the electron • Ernest Rutherford: Existence of the nucleus, and its relative size • Meitner & Fermi: Sustained nuclear fission • Ernest Lawrence: The cyclotron and trans-uranium elements

  4. Laws • Conservation of Mass • Law of Definite Proportion – • compounds have a constant composition. • They react in specific ratios by mass. • Multiple Proportions- • When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.

  5. Proof • Mercury has two oxides. • One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass. • Show that these compounds follow the law of multiple proportion. • Speculate on the formula of the two oxides.

  6. Dalton’s Atomic Theory (1808) • All matter is composed of extremely small particles called atoms • Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties John Dalton • Atoms cannot be subdivided, created, or destroyed • Atoms of different elements combine in simple whole-number ratios to form chemical compounds • In chemical reactions, atoms are combined, separated, or rearranged

  7. Modern Atomic Theory Several changes have been made to Dalton’s theory. Dalton said: Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Modern theory states: Atoms of an element have a characteristic average mass which is unique to that element.

  8. Modern Atomic Theory #2 Dalton said: Atoms cannot be subdivided, created, or destroyed Modern theory states: Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

  9. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

  10. Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

  11. Rutherford’s Gold Foil Experiment • Alpha particles are helium nuclei • Particles were fired at a thin sheet of gold foil • Particle hits on the detecting screen (film) are recorded

  12. Atomic Particles

  13. The Atomic Scale • Most of the mass of the atom is in the nucleus (protons and neutrons) • Electrons are found outside of the nucleus (the electron cloud) • Most of the volume of the atom is empty space “q” is a particle called a “quark”

  14. About Quarks… Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons”

  15. Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.

  16. Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011

  17. Molecules Two or more atoms of the same or different elements, covalently bonded together. Molecules are discrete structures, and their formulas represent each atom present in the molecule. Benzene, C6H6

  18. Covalent Network Substances Covalent network substances have covalently bonded atoms, but do not have discrete formulas. Why Not?? Graphite Diamond

  19. Ions • Cation: A positive ion • Mg2+, NH4+ • Anion: A negative ion • Cl-, SO42- • Ionic Bonding: Force of attraction between oppositely charged ions. • Ionic compounds form crystals, so their formulas are written empirically (lowest whole number ratio of ions).

  20. Periodic Table with Group Names

  21. The Properties of a Group: the Alkali Metals • Easily lose valence electron (Reducing agents) • React violently with water • Large hydration energy • React with halogens to form salts

  22. Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+

  23. Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Ba2+ Sr2+

  24. Predicting Ionic Charges Group 13: Loses 3 electrons to form 3+ ions B3+ Al3+ Ga3+

  25. Predicting Ionic Charges Group 14: Loses 4 electrons or gains 4 electrons Caution! C22- and C4- are both called carbide

  26. Predicting Ionic Charges Nitride N3- Group 15: Gains 3 electrons to form 3- ions P3- Phosphide As3- Arsenide

  27. Predicting Ionic Charges Oxide O2- Gains 2 electrons to form 2- ions Group 16: S2- Sulfide Se2- Selenide

  28. Predicting Ionic Charges F- Fluoride Br- Bromide Group 17: Gains 1 electron to form 1- ions Cl- Chloride I- Iodide

  29. Predicting Ionic Charges Group 18: Stable Noble gases do not form ions!

  30. Predicting Ionic Charges Many transition elements have more than one possible oxidation state. Groups 3 - 12: Iron(II) = Fe2+ Iron(III) = Fe3+

  31. Predicting Ionic Charges Some transition elements have only one possible oxidation state. Groups 3 - 12: Zinc = Zn2+ Silver = Ag+

  32. Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! ( ) 2. Check to see if charges are balanced. Ba2+ NO3- 2 Not balanced! 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

  33. Writing Ionic Compound Formulas Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES! ( ) NH4+ SO42- 2. Check to see if charges are balanced. 2 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!

  34. Writing Ionic Compound Formulas Example: Iron(III) chloride 1. Write the formulas for the cation and anion, including CHARGES! Fe3+ Cl- 2. Check to see if charges are balanced. 3 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!

  35. Writing Ionic Compound Formulas Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. Al3+ S2- 2 3 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!

  36. Writing Ionic Compound Formulas Example: Magnesium carbonate 1. Write the formulas for the cation and anion, including CHARGES! Mg2+ CO32- 2. Check to see if charges are balanced. They are balanced!

  37. Writing Ionic Compound Formulas Example: Zinc hydroxide 1. Write the formulas for the cation and anion, including CHARGES! ( ) 2. Check to see if charges are balanced. Zn2+ OH- 2 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!

  38. Writing Ionic Compound Formulas Example: Aluminum phosphate 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. Al3+ PO43- They ARE balanced!

  39. Naming Ionic Compounds • 1. Cation first, then anion • 2. Monatomic cation = name of the element • Ca2+ = calcium ion • 3. Monatomic anion = root + -ide • Cl- = chloride • CaCl2 = calcium chloride

  40. Naming Ionic Compounds(continued) Metals with multiple oxidation states • some metal forms more than one cation • use Roman numeral in name • PbCl2 • Pb2+ is the lead(II) cation • PbCl2 = lead(II) chloride

  41. Naming Binary Compounds • Compounds between two nonmetals • First element in the formula is named first. • Second element is named as if it were an anion. • Use prefixes • Only use mono on second element - P2O5 = diphosphorus pentoxide CO2 = carbon dioxide CO = carbon monoxide N2O = dinitrogen monoxide

  42. Acids • Substances that produce H+ ions when dissolved in water. • All acids begin with H. • Two types of acids: • Oxyacids • Non-oxyacids

  43. Naming acids • If the formula has oxygen in it • write the name of the anion, but change • ate to -ic acid • ite to -ous acid • Watch out for sulfuric and sulfurous • H2CrO4 • HMnO4 • HNO2

  44. Naming acids • If the acid doesn’t have oxygen • add the prefix hydro- • change the suffix -ide to -ic acid • HCl • H2S • HCN

  45. Formulas for acids • Hydrofluoric acid • Dichromic acid • Carbonic acid • Hydrophosphoric acid • Nitric acid • Perchloric acid • Phosphorous acid HF H2Cr2O7 H2CO3 H3P HNO3 HClO4 H3PO3

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