1. Ch. 4: Atomic Structure 4.1 Defining the Atom

2. History  Democritus named the most basic particle named the most basic particle atom- means “indivisible” atom- means “indivisible”  Aristotle didn’t believe in atoms didn’t believe in atoms thought matter was continuous thought matter was continuous

3. History  by 1700s, all chemists agreed: on the existence of atoms on the existence of atoms that atoms combined to make compounds that atoms combined to make compounds  Still did not agree on whether elements combined in the same ratio when making a compound

4. Dalton’s Atomic Theory-1803 1. Matter is made of small, indivisible particles called atoms 2. Atoms of same element have the same size, mass, and properties 3. Atoms of different element combine in whole number ratios to make compounds. 4. In chemical reactions, atoms are combined, separated, and rearranged.

5. Consequences of Dalton’s Theory  The “billiard ball model”-the atom is viewed as a small solid indivisible sphere. is viewed as a small solid indivisible sphere.  Some parts of Dalton’s theory were wrong: atoms are divisible into smaller particles (subatomic particles) atoms are divisible into smaller particles (subatomic particles) atoms of the same element can have different masses (isotopes) atoms of the same element can have different masses (isotopes)  Most important parts of atomic theory: all matter is made of atoms all matter is made of atoms atoms of different elements have different properties atoms of different elements have different properties

6. Law of Conservation of Mass  mass is neither created or destroyed during regular chemical or physical changes

7. Law of Definite Proportions  any amount of a compound contains the same element in the same proportions by mass No matter where the copper carbonate is used, it still has the same composition

8. Law of Multiple Proportions  applies when 2 or more elements combine to make more than one type of compound  the mass ratios of the second element simplify to small whole numbers

9. Ch. 4: Atomic Structure 4.2 Structure of Atom Discovery of the subatomic particles

10. Discovery of Electron  resulted from scientists (JJ Thompson) passing electric current through gases to test conductivity  used cathode-ray tubes  noticed that when current was passed through a tube a glow (or “ray”) was produced

11. Discovery of Electron Noted Qualities of Ray Produced: 1. existed- there was a shadow on the glass when an object was placed inside 2. had mass- the paddle wheel placed inside, moved from one end to the other so something must have been “pushing” it

12. Discovery of Electron Noted Qualities of Ray Produced: 3. negatively charged- the rays were attracted to the positive pole (anode)- opposites attract!

13. Discovery of Electron  Conclusion: there were negatively charged particles inside the cathode ray there were negatively charged particles inside the cathode ray  The particle was called ELECTRON  Cathode rays are made of electrons

14. Discovery of Electron  J.J. Thomson (English 1897) did more experiments to actually make the discovery  he found ratio of charge to mass of this particle  since the ratio stayed constant for any metal that contained it, it must be the same in all of the metals

15. Plum Pudding Model (1897) Plum Pudding Model (1897)  proposed by Joseph John Thomson  Nobel Prize in physics in 1906  the atom was a sphere of positive electricity (which was diffuse) with negative particles (electrons) imbedded throughout

16. Charge and mass of electron 1916- Robert Millikan discovered the charge and the mass of the electron discovered the charge and the mass of the electron Electron has  a charge of -1  a mass of 1/ 1840 of the mass of Hydrogen atom (the smallest atom)

17. Are electrons the only particles?  since atoms are neutral, something must balance the negative charge  since an atom’s mass is so much larger than the mass of its electrons, there must be other matter inside an atom

18. Canal rays 1886 Goldstein  In the cathode tube experiment Goldstein noticed another rays traveling from anode to cathode  He named those rays canal rays and are made of positive particles called protons

19. Discovery of Nucleus  Rutherford discovered the nucleus by shooting alpha particles (have positive charge) at a very thin piece of gold foil  1911- Ernest Rutherford- the gold foil experiment  he predicted that the particles would go right through the foil at some small angle

20. Discovery of Nucleus

21.  some particles (1/8000) bounced back from the foil  this meant there must be a “powerful force” in the foil to hit particle back Predicted ResultsActual Results

22. Discovery of Nucleus Characteristics of “Powerful Force”: 1. dense- since it was strong enough to deflect particle 2. small- only 1/8000 hit the force dead on and bounced back 3. positively charged- since there was a repulsion between force and alpha particles 4. “Powerful Force”= Nucleus

23. The Nuclear Model- proposed by Ernest Rutherford  the atom is mostly empty space with a dense positively charged nucleus surrounded by negative electrons.  the atom is mostly empty space with a dense positively charged nucleus surrounded by negative electrons.  Rutherford received the Nobel Prize in chemistry in 1908 for his contributions into the structure of the atom.  Rutherford received the Nobel Prize in chemistry in 1908 for his contributions into the structure of the atom.

24. The discovery of Neutrons 1932- Chadwick  Neutrons have NO CHARGE  Neutrons have the mass almost equal to the proton

25. Structure of Atom  Nucleus: contains protons and neutrons contains protons and neutrons takes up very little space takes up very little space  Electron Cloud: contains electrons contains electrons takes up most of space takes up most of space

26. Subatomic Particles  includes all particles inside atom proton proton electron electron neutron neutron  charge on protons and electrons are equal but opposite  to make an atom neutral, need equal numbers of protons and electrons

27. Subatomic Particles  number of protons identifies the atom as a certain element  protons and neutrons are about same size  electrons are much smaller  nuclear force- when particles in the nucleus get very close, they have a strong attraction proton + proton proton + proton proton + neutron proton + neutron neutron + neutron neutron + neutron

28. The Subatomic Particles ParticleSymbolCharge Relative Mass Electrone-1/1840 Protonp++11 Neutron n 0 01

29. Ch. 4 Atomic Structure 4.3 Distinguishing among Atoms

30. Atomic structure Atomic number Z Z= atomic number= # protons= # electrons  Indicates the position of the element in the periodic table (the whole number by each element)  Ex: 6 C, 19 K, 1 H Mass number A A= # protons + # neutrons  To find A round up the decimal number for each element in the Periodic table  Ex: 12 C, 39 K, 1 H or C-12, K-39, H-1

31. How to calculate  A. electrons= atomic number Z  B. protons= atomic number Z  C. neutrons= mass number – atomic number n= A-Z n= A-Z

32. ISOTOPES  Members of the same family  Have the same chemical symbol number of protons number of electrons atomic number  Have different mass numbers numbers of neutrons  Ex: 12 C, 13 C, 14 C, 16 C- all have 6 electrons and 6 protons but 6, 7, 8, and 10 neutrons

33. Relative Atomic Mass  since masses of atoms are so small, it is more convenient to use relative atomic masses instead of real masses  to set up a scale, we have to pick one atom to be the standard  since 1961, the carbon-12 nuclide is the standard and is assigned a mass of exactly 12 amu

34. Relative Atomic Mass  atomic mass unit (amu)- one is exactly 1/12 th of the mass of a carbon-12 atom  mass of proton= 1.007276 amu  mass of neutron= 1.008665 amu  mass of electron= 0.0005486 amu

35. Relative Atomic Mass  the mass number (A) and the relative atomic mass are very close but not the same because relative atomic mass includes electrons relative atomic mass includes electrons the proton and neutron masses aren’t exactly 1 amu the proton and neutron masses aren’t exactly 1 amu

36. Average Atomic Mass  weighted relative atomic masses of the isotopes of each element  Is the decimal number found in each box for each element in the Periodic table

37. Average Atomic Mass  To calculate it we need to know: number of isotopes number of isotopes Mass of each isotope Mass of each isotope Percentage (relative abundance) of each isotope) Percentage (relative abundance) of each isotope)

38. Calculating Average Atomic Mass  Naturally occurring copper consists of: 69.71% copper-63 (62.929598 amu) 69.71% copper-63 (62.929598 amu) 30.83% copper-65 (64.927793 amu) 30.83% copper-65 (64.927793 amu) (0.6971 x 62.929598)+(0.3083 x 64.927793) =63.55 amu

39. Calculating Average Atomic Mass  An element has three main isotopes with the following percent occurances: #1: 19.99244 amu, 90.51% #1: 19.99244 amu, 90.51% #2: 20.99395 amu, 0.27% #2: 20.99395 amu, 0.27% #3: 21.99138 amu, 9.22% #3: 21.99138 amu, 9.22%  Find the average atomic mass and determine the element.

40. Calculating Average Atomic Mass Neon