Early Atomic Theory and Structure

1. Early Atomic Theory and Structure

2. Chapter 5—Early Theories • What is stuff made of? • What makes something move? • How do we know it’s alive? • Is there a fundamental particle that everything is made up of? • Is there a universalconstant to all matter?

3. Chapter 5.1 Early Thoughts • Roots of atomic theory are as old as 440 B.C. with Democritus’ idea of the atom • It took 2 000 years for us to expand on this idea. The new theory was to be done by an English schoolmaster John Dalton in the early 1800s.

4. Dalton’s Atomic Theory • His theory included 6 postulates • 1. Elements are made up of atoms • 2. Atoms of the same element are alike • 3. Atoms of different elements are different by virtue of their size and mass • 4. Chemical compounds are formed by the union of two or more atoms of different elements • 5. Atoms combine to form compounds in whole number ratios (1:2 or 2:2, etc] • 6. Atoms of two elements may combine in different ratios to form more than one compound

5. 5.2 Dalton’s Atomic Theory (cont.) • Why is it a Theory? • Which are still true? • Which do we know More info about now?

6. 5.3 Composition of Compounds • The Law of Definite Compositionstates that a compound always contains two or more elements combined in definite proportion by mass

7. Law of Multiple Proportions • The Law of Multiple Proportions states that atoms of two or more elements may combine in different ratios to produce more than one compound

8. 5.4 - 5.8 Subatomic Particles • Through the years of the late 1800s and into the early 1900s it was determined that there are three subatomic particles • Electrons (discovered first) • Protons (reasoned to exist if elements are neutral) • Neutrons (discovered last)

9. Electrons • 1. Electron which occupiesthe areaoutsidethe nucleusand has anegativecharge,relative to the other subatomic particles it has negligible (so small that it can be ignored) mass.

10. Protons • 2. Protonwhich existsin the nucleus, has apositive chargeand has mass roughly equal to neutrons

11. Neutron • 3. Neutrons(discovered last] functions as the glue that holds the nucleus together so that the protons don’t repel each other, it hasno chargeand roughly the same mass as the proton

12. Isotopes • Isotopeshave same number of protons (so they are the same element) but different number of neutrons • Some isotopes are radioactive

13. Atomic Number • Atomic Number= the number of protons; unique to each element and the way the periodic table is arranged

14. Mass number • Mass Number= protons + neutrons (whole number • Cannot be found on the periodic table!

15. Check yourself • The nucleus is made up of what two types of subatomic particles?

16. Formulas you should know • Atomic number = # of protons • In an atom (uncharged): • # of protons = # of electrons • Mass # = # protons + # neutrons or • # neutrons = mass # - # protons • Charge = # protons - # electrons (for ions) • Remember the atomic # and # of protons give the element its identity and does not change

17. Elements composed of atoms Elements or atoms in an unbonded state have the same number of electrons as protons (They are neutral)

18. Ions Ions have an unequal number of electrons and protons. An atom loses or gains electrons to take on a charge (protons/neutrons are not transferred) Charge = #protons - # electrons

19. Ionic Charge • Charge is written in the upper RIGHT corner of the element’s symbol. • It is written with the number first and the sign second unless it is a + 1 or a -1 in which case it is just written as + or -. • Negative ions change their names to end in –ide like fluorine is fluoride

20. Ionic Notation X3- This means that this element has a -3 charge.

21. Self Check • What is the charge of a substance with 14 protons, 15 neutrons, and 14 electrons?

22. Self Checker • If a substance has a charge of +2, this means that the number of protons is (circle one: LESS than or GREATER than) the number of electrons?

23. ISOTOPIC NOTATIONisotopes are atoms with the same number of protons but different number of neutrons A Z X A = mass number (the total number of protons + neutrons) Z = atomic number (the total number of protons) X = element symbol

24. READING ISOTOPIC NOTATION 46 21 Sc 46 = mass number (the total number of protons (21) + neutrons (25) 21 = atomic number (the total number of protons (21)) Sc = element symbol In a neutral atom, the number of electrons (21) is equal to the number of protons.

25. PRACTICE PROBLEMS 7 8 15N # protons = ____ # neutrons= ____ #electrons = ___ 35P # p = ____ # n= ____ #e- = ___ 62Cu2+ # p = ____ # n= ____ #e- = ___ 76Se3- # p = ____ # n= ____ #e- = ___ 7 15 15 20 29 33 27 42 34 37

26. Writing ISOTOPIC NOTATION • Write the symbol for the atom with an atomic number of 21 and a mass number of 48. • Give the complete chemical notation for the nuclide with 23 protons, 26 neutrons and 20 electrons. • Write the isotopic notation for • Z = 46 A = 110 • An atom containing 24 protons, 28 neutrons, and 21 electrons • Titanium-50 48 Sc 49V3+ 110Pd 52Cr3+ 50Ti

27. PRACTICE PROBLEMS • 196 Pt4+ # p = _____ # n = _____ #e- = _____ mass number = ________ atomic number = _______ atomic mass = ________ name of element = _______ 2. Indicate the appropriate atomic mass of an element with 30 protons, 30 neutrons, and 28 electrons. 74 118 78 78 196 195.1 amu platinum 65.39 amu

28. Atomic Mass • Atomic Mass= number on the periodic table reflecting the mass all isotopes known and their relative percentages (on periodic table below element’s symbol--usually not a whole number)

29. Atomic Mass • The atomic mass of an element represents the average mass of all the isotopes found in nature. No element exists with only one possible isotope. Hydrogen has the smallest number of isotopes: 1H protium, 2H deuterium, 3H tritium. Its atomic mass is 1.0079 amu (atomic mass units). The atomic mass is calculated by adding the % of 1H mass found in nature to the % of 2H mass found in nature plus the % of 3H mass. • % 1H + % 2H + % 3H = average mass (atomic mass) • Generally the formula used is: % X + % Y + % Z… = atomic mass. An instrument called the mass spectrometer is generally used to determine the percentages and individual masses of each isotope.

30. Atomic Mass • Silver is found to have two stable isotopes, one has an atomic mass of 106.904 amu and the other weighs 108.905 amu. The first isotope represents 51.82 % of the mass of the element and the second represents 48.18 %. What is the atomic mass of the element silver? The equation to use is %X + % Y = average And remember to turn your percents into fractions before multiplying. (0.5182) 106.904 amu + (0.4818) 108.905 amu =? 55.398 amu + 52.470 amu =? 107.868 amu !! Now look at the periodic table to verify the answer.

31. PRACTICE PROBLEMS # 8 1. A sample of neon contains three isotopes, neon-20 (with an isotopic mass of 19.9924 amu), neon-21 (20.9939 amu) and neon-22 (21.9914 amu). The natural abundances of these isotopes are 90.92%, 0.257 %, and 8.82 %. Calculate the atomic weight of neon. 2. There are only two naturally occuring isotopes of copper, 63Cu and 65Cu. Copper has an atomic mass of 63.55 amu. What is the natural abundance of each isotope? 3. There are only two naturally occuring isotopes of gallium, 69Ga and 71Ga. What is the natural abundance of each isotope? 20.17 amu 65Cu = 30% & 63Cu = 70% 69Ga = 60% and 71Ga = 40%

32. GROUP STUDY PROBLEM #8 _______1. The element with atomic number 53 contains a) 53 neutrons b) 53 protons C) 26 neutrons & 27 protons d) 26 protons & 27 neutrons _______2. The mass of one atom of an isotope is 9.746 x 10-23 g. One atomic mass unit has the mass of 1.6606 x 10-24 g. The atomic mass of this isotope is a) 5.870 amu b) 16.18 amu c) 58.69 amu d) 1.627 amu 108 _______3. The number of neutrons in an atom of 47 Ag is a) 47 b) 108 c) 155 d) 61 27 _______4. The number of electrons in an ion of 13 Al3+ is a) 13 b) 10 c) 27 d) 14 _______5. What is the relative atomic mass of boron if two stable isotopes of boron have the following mass and abundance: 10.0129 amu (19.91%) & 11.0129 (80.09%) a) 10.81 amu b) 10.21 amu c) 10.62 amu d) 10.51 amu

33. Test your Knowledge

34. Test your Knowledge

35. Table Information Symbol H Atomic Number 1 Atomic Weight 1.00794 Oxidation States +1, -1 Electronegativity, Pauling 2.2 State at RT Gas, Non-metal Melting Point, K 14.01 Boiling Point, K 20.28 Hydrogen

36. The Periodic Table • Horizontal rows are called periods • Vertical columns are called groups • We will use 1- 18 as group designations. • Group 1 is Alkali Metals • Group 2 is the Alkaline Earth Metals • Group 18 Inert or Noble Gases • Group 17 Halogens

37. Larger Groups • Groups 3 –12 are the heavy metals or transition elements • Two periods at the bottom are called the rare earth elements or the inner transition elements.