1. Chapter 3 ATOMS: THE BUILDING BLOCKS OF MATTER

2. Section 1 FROM PHILOSOPHICAL IDEA TO SCIENTIFIC THEORY

3. Foundations of Atomic Theory  Particle Theory of Matter  Democritus in 400 B.C.  Stated that nature’s basic particle was the atom (“indivisible” in Greek).  Aristotle  Believed all matter was continuous (could be divided forever), and did not believe in atoms.  Neither had experimental evidence to support their claims.

4. Fast Forward to the18 th century  Three laws discovered due to improved instrumentation and carefully observed chemical reactions  Law of conservation of mass—Mass is neither created nor destroyed during an ordinary chemical reaction or physical change.  Law of definite proportions—A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.  Example: Table salt will always consist of 39.34 % Na and 60.66% Cl.  Law of multiple proportions—If 2 or more compounds are composed of the same two elements, the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.  Example: Carbon and Oxygen can combine to form CO 2 or CO.

5. Practice (law of conservation of mass)  If 3.5 g of X reacts with 10.5 g of Y to form the compound XY, what is the percent by mass of X in the compound?  If 40 g of X reacts with 35 g of Y, what is the mass of the product XY?  If 3.5 g of X reacts with 10.5 g of Y to form the compound XY, how many grams of Y would react to form XY 2 ? What would be the final mass of XY 2 ?

6. Practice (law of definite proportions)  2 unknown compounds are tested. Compound 1 contains 15.0g of hydrogen and 120.0g oxygen. Compound 2 contains 2.0g of hydrogen and 32.0g oxygen. Are the compounds the same?

7. Practice with Multiple Proportions  Three compounds containing K and O are compared. Analysis shows that for each 1.00 g of O, the compounds have 1.22 g, 2.44 g, and 4.89 g of K, respectively. Show how these data support the law of multiple proportions.  First: Find ratios of K for the different compounds.  If the ratios are whole numbers, it supports the law.

8. More practice with Multiple proportions…  In 100 g of the compound A there are 57.1 g O and 42.9 g C. In 100 g of the compound B, there are 72.7 g O and 27.3 g C. Show how this data supports the law of multiple proportions.

9. Recall  1.) If 13 g of X reacts with 45 g of Y to form the compound XY, what is the percent by mass of X in the compound?  2.) Two unknown compounds are tested. Compound 1 contains 32.6 g of hydrogen and 167.4 g of Carbon. Compound 2 contains 8.0 g of hydrogen and 24.0 g of carbon. Are the compounds the same?

10. Dalton’s Atomic Theory  John Dalton—proposed Atomic Theory  Accounted for all three laws: conservation of mass, definite proportions, and multiple proportions.  Dalton’s Atomic Theory consists of the following statements:  All matter is composed of atoms.  Atoms of a given element are identical in size, mass, and all other properties.  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 separated, combined, or rearranged.

11. Modern Atomic Theory  Advances in instrumentation have allowed some aspects of Dalton’s theory to be proven incorrect.  Example: The thoughts that atoms are not divisible into smaller particles and have the exact same mass are incorrect.  The two most important aspects of Dalton’s theory still hold true  All matter is composed of atoms.  Atoms of any one element differ in properties than atoms of another element.

12. Section 2 THE STRUCTURE OF AN ATOM

13. Discovery of the Electron  Scientific advances allowed for the discovery of smaller particles that made up atoms.  William Crookes—Passed and electric current through a cathode ray tube.  Noticed that current passing through the tube produced a stream of glowing particles (particles were traveling from the negative end of the tube (cathode) to the positively charged end (anode).

14. Two conclusions resulted from experiments: 1.) Cathode rays were deflected by a magnetic field in the same manner as wire carrying an electric current (negatively charged). 2.) The rays deflected away from a negatively charged object.

16. Charge of Cathode Ray

17. JJ Thomson, 1900  Proved Dalton’s solid atom wrong when he discovered the particles had a mass and it was much less than the mass of a Hydrogen atom.  Concluded that the electron has a very large charge-to-mass ratio  Conducted Cathode-ray experiments to prove that atoms are divisible.  Proposed plum pudding model for the atom.

18. Milikan and the oil drop experiment Meaured mass of droplets Applied charge to droplets through x-rays. Applied electric voltage to top and bottom plates Measured the amount of voltage it took to keep the oil droplets suspended. Was able to measure the charge of a single electron! http://www.youtube.com/watch?v=EV1owO1H2dAwww.youtube.com/watch?v=EV1owO1H2dA

19. Ernest Rutherford  Discovered the nucleus of an atom with his gold-foil experiment.  http://www.youtube.com/watch?v=XBqHkraf8iE http://www.youtube.com/watch?v=XBqHkraf8iE  Concluded that the volume of the nucleus was very small compared to the total volume of the atom.

20. Sizes of Atoms  Expressed in picometers (pm).  1 pm = 1*10 -12 m = 0.000000000001m

21. Section 3 COUNTING ATOMS

22. Atomic Number The number of protons in each atom of a given element Unique to each element and does not change Used to identify elements.

23. Practice 1. How many protons and electrons are in each atom?  Radon, Rn  Titanium, Ti 2. An atom of an element contains 66 electrons. Which element is it? 3. An atom of an element contains 14 protons. Which element is it?

24. Isotopes and Mass Number  The amount of protons and electrons are constant in all neutral atoms of an element, but the number of neutrons can vary  Isotopes– atoms of the same element that have different numbers of neutrons  Example: Hydrogen atoms have three different isotopes.  Protium—1 proton/0 neutrons (accounts for 99.985% of all hydrogen atoms).  Deuterium—1 proton/1 neutrons (0.015%)  Tritium—1 proton/2 neutrons (radioactive, and can be produced artificially)  Most elements contain mixtures of isotopes. Tin (Sn) has 10 stable isotopes, the most of any element.  Mass Number-- The number of protons and neutrons that make up the nucleus of an isotope. Since electrons are so small, their mass is considered insignificant.

25. Designating Isotopes Isotopes are usually identified by their mass number.  Two Methods:  Hyphen notation—has hyphen after name of element, followed by the mass number.  Example: Tritium is hydrogen-3 since it has 2 neutrons and 1 proton.  Example: Uranium-235  Nuclear symbol  Mass number written as superscript  Atomic number written as subscript  Followed by element symbol

26. In both methods…  The number of neutrons is found by subtracting the atomic number from the mass number.  Example: Uranium-235 or 235 92 U  (Mass #) – (Atomic #) = (# of neutrons)  235 (protons + neutrons) – 92 protons= 143 neutrons

27. Notation Practice

28. Pop quiz

29. Relative Atomic Masses  Masses of subatomic particles are extremely small and hard to work with—scientists created a new unit  Atomic mass units (amu) – based on a standard of carbon-12 that has a mass of 12 amu  Neutron – 1.008665 amu = 1.66x10 -24 g  Proton – 1.007276 amu  Electron – 0.0005486 amu

30. Average Atomic Mass  Avg. atomic mass – the weighted average of the masses of the isotopes of that element  Avg. atomic mass = (mass x abundance) isotope1 + (mass x abundance) isotope2 + …  To find abundance – divide the percentage by 100 (all abundances must be decimals)  Atomic mass can help you determine which isotopes of that element is the most abundant.

31. Atomic Mass Practice 1. Boron has two naturally occurring isotopes: boron-10 (abundance 19.8%, mass 10.013 amu) and boron-11 (abundance 80.2%, mass 11.009 amu). Calculate the atomic mass of boron. 2. Nitrogen has two naturally occurring isotopes, N-14 and N-15. Its atomic mass is 14.007. Which isotope is more abundant? Explain.

32. Calculate the average atomic mass of lithium, which occurs as two isotopes that have the following atomic masses and abundances in nature: 6.017 amu, 7.30% and 7.018 amu, 92.70%.

33.  What is the atomic mass of hafnium if, out of every 100 atoms, 5 have a mass of 176, 19 have a mass of 177, 27 have a mass of 178, 14 have a mass of 179, and 35 have a mass of 180.0?  Iodine is 80% 127 I, 17% 126 I, and 3% 128 I. Calculate the average atomic mass of iodine.

34.  Magnesium has three naturally occuring isotopes. 78.70% of Magnesium atoms exist as Magnesium-24 (23.9850 g/mol), 10.03% exist as Magnesium-25 (24.9858 g/mol) and 11.17% exist as Magnesium-26 (25.9826 g/mol). What is the average atomic mass of Magnesium?

35. Relating Mass to Number of Atoms  The Mole (mol)  Since atoms are so tiny, we count them in groups  Same as 1 dozen = 12 donuts  Mole = SI unit for amount of substance.  The amount of substance that contains as many particles as there are atoms in exactly 12 g of carbon-12.  Avogadro’s number  The number of particles in exactly 1 mol of a pure substance.  6.022 × 10 23 atoms in 1 mol of any substance.

36. Practice  How many atoms are in 5.76 moles of a pure substance?  How many moles are in 8.56 × 10 26 atoms?

37. Molar Mass  The mass (in grams) of 1 mol of a pure substance.  Units are g/mol  Equal to the sum of the atomic masses of each element of a compound.  Example:  K 2 CO 3  2(39.10 g/mol) + 1(12.01 g/mol) + 3(16.00 g/mol) = 138.21 g/mol K CO

38. Gram/Mole Conversions  Use molar mass to convert moles to grams and vice versa.  Practice:  Convert 7.00 g of He to moles.  Convert 5.67 mol of Xe to grams.  Convert 65.4 g of K 2 CO 3 to mol.  Convert 5.21 mol of NaCl to grams.

39. Recall  Find the molar masses of the following compounds:  SrCO 3  C 13 H 18 O 2  Sn 2 (SO 4 ) 4  A sample of caffeine contains 96.1 g of carbon for every 10.1 g of hydrogen. If another sample of caffeine contains 30.0 g of carbon, how many g of hydrogen does it contain?

40. More Practice  How many atoms are in 4.56 g of Na 2 CO 3 ?  What is the mass of 5.67 × 10 32 atoms of NaOH?

41. Chapter 3 Vocabulary  Law of Conservation of Mass  Law of Definite Proportions  Law of Multiple Proportions  Atom  Nuclear Forces  Atomic Number  Isotope  Mass Number  Nuclide  Atomic Mass Unit  Average Atomic Mass  Mole  Avogadro’s number  Molar Mass