2. Atoms, Molecules, and Ions

3. Chemistry Timeline #1 B.C. 400 B.C. Demokritos and Leucippos use the term "atomos” 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  2000 years of Alchemy

4. 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

5. 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.

6. 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.

7. Dalton’s Atomic Theory (1808)  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  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

8. 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.

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

10. 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.

11. 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.

12. 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

13. Atomic Particles ParticleChargeMass (kg)Location Electron9.109 x 10 -31 Electron cloud Proton+11.673 x 10 -27 Nucleus Neutron01.675 x 10 -27 Nucleus

14. 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”

15. 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”

16. Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. IsotopeProtonsElectronsNeutronsNucleus Hydrogen–1 (protium) 110 Hydrogen-2 (deuterium) 111 Hydrogen-3 (tritium) 112

17. Atomic Masses IsotopeSymbolComposition of the nucleus % in nature Carbon-12 12 C6 protons 6 neutrons 98.89% Carbon-13 13 C6 protons 7 neutrons 1.11% Carbon-14 14 C6 protons 8 neutrons <0.01% Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011

18. 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, C 6 H 6

19. Covalent Network Substances Covalent network substances have covalently bonded atoms, but do not have discrete formulas. Why Not?? GraphiteDiamond

20. IonsIons  Cation: A positive ion Mg 2+, NH 4 +Mg 2+, NH 4 +  Anion: A negative ion  Cl , SO 4 2   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).

21. Periodic Table with Group Names

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

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

24. Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+

25. Predicting Ionic Charges Group 13: Loses 3 Loses 3 electrons to form 3+ ions B 3+ Al 3+ Ga 3+

26. Predicting Ionic Charges Group 14: Loses 4 Loses 4 electrons or gains 4 electrons Caution! C 2 2- and C 4- are both called carbide

27. Predicting Ionic Charges Group 15: Gains 3 Gains 3 electrons to form 3- ions N 3- P 3- As 3- Nitride Phosphide Arsenide

28. Predicting Ionic Charges Group 16: Gains 2 Gains 2 electrons to form 2- ions O 2- S 2- Se 2- Oxide Sulfide Selenide

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

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

31. Predicting Ionic Charges Groups 3 - 12: Many transition elements Many transition elements have more than one possible oxidation state. have more than one possible oxidation state. Iron(II) = Fe 2+ Iron(III) = Fe 3+

32. Predicting Ionic Charges Groups 3 - 12: Some transition elements Some transition elements have only one possible oxidation state. have only one possible oxidation state. Zinc = Zn 2+ Silver = Ag +

33. Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! Ba 2+ NO 3 - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2

34. Writing Ionic Compound Formulas Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES! NH 4 + SO 4 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2

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

36. Writing Ionic Compound Formulas Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ S 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! 23

37. Writing Ionic Compound Formulas Example: Magnesium carbonate 1. Write the formulas for the cation and anion, including CHARGES! Mg 2+ CO 3 2- 2. Check to see if charges are balanced. They are balanced!

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

39. Writing Ionic Compound Formulas Example: Aluminum phosphate 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ PO 4 3- 2. Check to see if charges are balanced. They ARE balanced!

40. Naming Ionic Compounds 1. Cation first, then anion1. Cation first, then anion 2. Monatomic cation = name of the element2. Monatomic cation = name of the element Ca 2+ = calcium ionCa 2+ = calcium ion 3. Monatomic anion = root + -ide3. Monatomic anion = root + -ide Cl  = chlorideCl  = chloride CaCl 2 = calcium chlorideCaCl 2 = calcium chloride

41. Naming Ionic Compounds (continued)  some metal forms more than one cation  use Roman numeral in name PbCl 2 Pb 2+ is the lead(II) cation PbCl 2 = lead(II) chloride Metals with multiple oxidation states

42. 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 - P 2 O 5 = CO 2 = CO = N 2 O = diphosphorus pentoxide carbon dioxide carbon monoxide dinitrogen monoxide

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

44. 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 H 2 CrO 4 HMnO 4 HNO 2

45. Naming acids If the acid doesn’t have oxygen add the prefix hydro- change the suffix -ide to -ic acid HCl H 2 S HCN

46. Formulas for acids Hydrofluoric acid Dichromic acid Carbonic acid Hydrophosphoric acid Nitric acid Perchloric acid Phosphorous acid HF H 2 Cr 2 O 7 H 2 CO 3 H3PH3P HNO 3 HClO 4 H 3 PO 3