Atomic Structure

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Atomic Structure

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1. Atomic Structure

2. Standards • 1a. Students know how to relate the position of an element in the periodic table to its atomic number and atomic mass. • 1b. Students know how to use the periodic table to identify metals, semi-metals {metalloids}, non-metals, and halogens • 1e. Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass.

3. Purpose • We will use this information to build our chemistry knowledge. • We will use this information as the foundation to calculate limiting reagent problems. • The standardized exams in the spring will test you on this information.

4. Objectives • Know the 3 particles of the atom and where they reside • Know the difference between atomic number and mass number • Know how to write nuclide symbols • Know the three isotopes of hydrogen • Know how to calculate atomic mass • Know how to calculate percent composition

5. First Some Questions… • What are atoms made up of? • Protons, Neutrons and Electrons • Where do you find these particles? • Protons and Neutrons are located in the nucleus • Electrons are located in the outer rings, outside the nucleus.

6. Vocabulary • Atom- from the Greek atomos=indivisible. The atom is the smallest particle of an element that retains the properties of that element. • Nucleus: the center of the atom; composed of neutrons and protons. Because the mass of the proton and the neutron is much larger than that of electrons, almost all the mass is located in the nucleus. • Ion: a charged particle; # protons ≠ # electrons • Electrons occupy most of the volume of an atom outside/around the nucleus.

7. Fundamental Particles • Proton • A positively charged particle located in the nucleus. • Neutron • A neutral particle located in the nucleus. • Electron • A negatively charged particle located outside the nucleus.

8. Question • What differentiates one atom from another atom? • The number of PROTONS

9. Atomic Number (Z) • Number of protons in the nucleus of an atom • This number is found on the Periodic Table • Atomic Number identifies an element • Always a positive number (b/c it is a counting #) • Tells number of electrons in a neutral atom • An atom is electrically neutral

10. What does it mean to be electrically neutral? • The atom has no charge • The number of protons = the number of electrons

11. Question • What observations can you make about atomic numbers on the periodic table? • Atomic Number increases as you go across the rows from left to right.

12. Questions • What is the atomic number of Chlorine? • What can you tell me about its protons and electrons? • What element has 20 protons? • What is the relationship between the # protons and the atomic number? • They’re equal.

13. Complete the Chart K 19 19 Boron B 5 Sulfur S 16 Yttrium 39 39

14. Mass Number (A) • Total number of protons and neutrons in the nucleus of an atom • Always a positive number • You can determine the nuclear composition of an atom from its mass number and atomic number

15. Question • What do the atomic number and the mass number have in common? • Both Positive integers • Both havethe same # of protons

16. How to find # of Neutrons • Mass # - Atomic#= # Neutrons • Or • # protons + # neutrons= Mass # • (atomic number + # neutrons)=Mass #

17. Complete the Chart 19 9 9 F 14 29 14 Si 22 22 25 Ti 25 30 25 Mn 6 12 6 6

18. Isotopes • Atoms of the same element with differing numbers of neutrons • Atoms with the same atomic number but different mass number • Isotopes of an element have different masses • Chemical properties of different isotopes are virtually the same

19. Nuclide Symbol • A=Mass # • Z= Atomic #

20. Nuclide • A specific kind of atom • Specification of an element in terms of its nuclear composition/structure • Tells number of protons and number of neutrons

21. Complete the Chart 19 9 9 F 14 29 14 Si 22 22 25 Ti 25 30 25 Mn 6 12 6 6

22. 3 Isotopes of Hydrogen Protium 1 0 1 Deuterium 1 1 1 Tritium 1 2 1

23. Nuclides • By specifying the nuclear structure, then you call it a nuclide. • But if you say Carbon atom, you do not know which Carbon atom it is, therefore you don’t know how many neutrons it has • Example: Brothers and Sisters- • You are members of the Jones family, but you have not specified which Jones member you are referring to.

24. Write the nuclide name and nuclide symbol Chlorine-37 Anion Calcium-40 Cation Uranium-238

25. Atomic Mass • A weighted average of the atoms in a naturally occurring sample of the element. • Naturally occurring: no matter where you get the sample from, it will have the same percentages of isotopes.

26. Construct a Fruit Basket • Fruit TypeWeight of Each Piece 2 grapefruit 14 oz 4 apples 10 oz 3 pears 7 oz 1 kiwi 3 oz • What is the Average Weight?

27. Fruit Basket • Average weight=9.2oz • Each type of fruit makes a different contribution to the overall weight • How many pieces of fruit actually weigh 9.2 ounces? • None! • What does 9.2 oz mean? • Fictitious non-existent piece of fruit

28. Atomic Mass • If you have a recipe, you could count items to put in, say 200 chocolate chips, 3 eggs, etc. • But suppose I have a recipe to make a compound. • I need 100 hydrogen atoms and 50 oxygen atoms-you cannot count atoms or pluck them out with atomic tweezers! • So instead they mass them (weigh them) • Careful here, the mass of an object is completley different from the weight of an object.

29. Question • What accounts for the mass of the atom? • # protons & # neutrons in the nucleus

30. Atomic Mass • Know that 1.0 amu is defined as exactly 1/12 the mass of a atom. • Carbon-12 has 6 protons and 6 neutrons, therefore 1 proton or 1 neutron = ~1 amu • 1 amu = 1.6606 x 10 -24 grams • Since the mass mostly depends on # protons and # neutrons, you’d think atomic mass would be a whole number, but it isn’t. How come?

31. Atomic Mass • In nature, most elements exist as a mixture of 2 or more isotopes. • Each isotope of an element has a fixed, constant mass and fixed constant relative abundance. • Relative abundance- (amount) X 100=% (how much of each isotope is present) • Sample of carbon from anywhere in the world; coal from S. Africa, W. Virginia or Pennsylvania → 99% C-12 and 1% C-13 • Atomic Mass of periodic table takes into account the larger and smaller masses of the isotopes, just like the average piece of fruit accommodated the larger and smaller masses. • → Idea of weighted average

32. Calculating Atomic Mass • To calculate atomic mass you need to know 3 things: • # of stable isotopes • Mass of each isotope • % abundance of each isotope • Each isotope is a piece of fruit and the isotope’s mass is the weight of each piece of fruit.

33. Example: Chlorine Calculation • mass of isotope X relative abundance + mass of isotope X relative abundance =_______amu • (34.969)(.7553) + (36.935)(.2447) = • That’s the same value on the periodic table! 35.4500amu

34. Question • How many chlorine atoms actually have a mass of 35.45 amu? • NONE • So the atomic mass, in amu, is the average of a fictitious non-existent atom of an element.

35. Example: Copper Calculation (62.9298)(.6909)+(64.9278)(.3091)= 63.5464 amu

36. Calculating Relative Abundance • To Calculate % Abundance: • Make a Chart • Isotopic Mass X %Abundance of each isotope • Set-up equation • Solve for “x” • Plug in “x” value to solve for “y”

37. Example x 1- x 1.00 x + y = 1.00 y = 1 – x 10.103 (x) + 11.009 (1 –x) = 10.811 10.103x + 11.009 -11.009x = 10.811 -0.996x = -0.198 x = .1987 y= 1-.1987 y= .8013 B-10 = 19.87% B-11 = 80.13%

38. The End