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Atoms, Molecules, and Ions

Atoms, Molecules, and Ions

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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. 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

  5. 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.

  6. 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!

  7. 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.

  8. 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.

  9. Mass of the Electron 1909 – Robert Millikan determines the mass of the electron. The oil drop apparatus Mass of the electron is 9.109 x 10-31 kg

  10. Conclusions from the Study of the Electron • Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. • Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons • Electrons have so little mass that atoms must contain other particles that account for most of the mass

  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. Try it Yourself! In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?

  13. The Answers Target #1 Target #2

  14. Rutherford’s Findings • Most of the particles passed right through • A few particles were deflected • VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions: • The nucleus is small • The nucleus is dense • The nucleus is positively charged

  15. Atomic Particles

  16. 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 ofthe atom is empty space “q” is a particle called a “quark”

  17. 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”

  18. Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.

  19. Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p+ + n0 18 8 8 18 Arsenic 75 33 75 Phosphorus 16 15 31

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

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

  22. 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

  23. Periodic Table with Group Names

  24. 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 theirformulas are writtenempirically (lowest whole number ratio of ions).

  25. 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

  26. Predicting Ionic Charges Group IA: Lose 1 electron to form1+ions H+ Li+ Na+ K+

  27. Predicting Ionic Charges Group IIA: Loses 2 electrons to form2+ions Be2+ Mg2+ Ca2+ Ba2+ Sr2+

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

  29. Predicting Ionic Charges Group IVA: loses 4 electrons or gains 4 electrons

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

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

  32. Predicting Ionic Charges F1- Fluoride Br1- Bromide Group VIIA: Gains 1 electron to form 1-ions Cl1- Chloride I1- Iodide

  33. Predicting Ionic Charges Group VIII: StableNoble gasesdo notform ions!

  34. Predicting Ionic Charges Groups 3 – 12: Manytransitionelements have more than one possible oxidation state. Groups 3 - 12: Iron(II) = Fe2+ Iron(III) = Fe3+

  35. Predicting Ionic Charges Groups 3 – 12:Sometransitionelements have only one possible oxidation state. Zinc = Zn2+ Silver = Ag+

  36. 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, usingsubscripts. Use parenthesesif youneed more than oneof apolyatomic ion.

  37. 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!

  38. 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,usingsubscripts. Use parenthesesif youneed more than oneof apolyatomic ion. Not balanced!

  39. 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,usingsubscripts. Use parenthesesif youneed more than one of a polyatomic ion. Not balanced!

  40. 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!

  41. 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!

  42. 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!

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

  44. Naming Ionic Compounds(continued) Metals with multiple oxidation states • some metal forms more than onecation • useRoman numeralin name PbCl2 • Pb2+is the lead(II) cation • PbCl2 = lead(II) chloride

  45. Naming Binary Compounds • Compounds between twononmetals • First element in the formula isnamed first. • Second element is named as if it were ananion. • Use prefixes • Only usemonoon second element - P2O5 = diphosphoruspentoxide CO2 = carbondioxide CO = carbonmonoxide N2O = dinitrogenmonoxide