1. Review for unit 2 test

2. How do elements differ from each other?  Most obviously in their properties (reactivity, m.p., b.p., solubility, density, relative weight.)  They also differ in their identity/makeup, but that is easiest to consider at the atomic level.

3. What unit would be used to discuss the mass of an atom?  a.m.u. (atomic mass units).  Example: A copper atom has a mass of 63.5 amu  A lead atom has a mass of 207.2 amu.

4. We did a worksheet comparing their weights:  Originally, every element was weighed in terms of the lightest element (hydrogen)  Helium has a mass 4 times that of hydrogen, thus helium has a mass of 4 amu.  Lithium has a mass 7 times that of hydrogen; lithium’s mass is 7 amu.

5. Today the relative mass is not based on hydrogen.  The reason is simply in order to be more exact.  The weights are based off of (compared to) the carbon-12 isotope.

6. Because atoms are so small, and have such little mass, we introduced the mole concept. What is a mole?  Simply stated, it’s just a quantity  (a number)

7. What is special about the mole quantity?  It’s the number of particles it takes to get the same mass in grams as what one particle’s mass is in amu’s.

8. Examples:  1 sodium (Na) atom is 22.99 amus.  To get 22.99 grams of sodium, you need 1 moles worth of sodium atoms.  So, a mole is the quantity of sodium atoms needed to get 22.99 grams

9. Examples:  1 mercury (Hg) atom is 200.59 amus.  To get 200.59 grams of mercury, you need 1 moles worth of mercury atoms.  So, a mole is the quantity of mercury atoms needed to get 200.59 grams of mercury

10. Why is the mole concept useful to us?  It allows us to “count by weighing”.

11. Example  If you have 5 lbs of nails, how many nails do you have?  You first need to know the mass of one nail.  If the mass of one nail is 0.02857 lbs  5 lbs/0.02857 = 175 nails.

12. We do the same counting method using moles.  If you have 200. grams of iron, how many iron atoms do you have?  A mole of iron is 55.845 grams  200/55.845 = 3.58 moles of iron  This would be (3.58)*(6.022x10 23 ) = 2.16 x 10 24 iron atoms

13. What are three basic parts of the periodic table?  Metals (most elements are metals.  Nonmetals.  A handful of “semi-metals” or “metaliods.

14. The Periodic table

15. Why is it called the “periodic table”?  Because it has a “periodic” (repeating) quality to it.  The repeating patterns are seen within the group/family (vertical columns)  So: Li, Na, K, Rb, Cs all have similar properties-hence the term the “alkali metal” family.

16. Who arranged the “periodic table”?  Dmetri Mendelev (1869)

17. How did Mendelev arrange his table of elements?  In order of generally increasing atomic weight (with a few exceptions)  With repeating properties—most notably the property of “valence”.  Also considered were properties such as density, melting point, boiling point, atomic weight

18. What is valency?  It’s the ratio in which atoms combine-particularly with other elements…especially hydrogen, oxygen, and chlorine  Example: Nitrogen combines with hydrogen in a 1:3 ratio, so we say nitrogen has a “valency” of 3.

19. Given the following relationship, what is the “valency” of the elements below?  4.0 grams of element X combines with 39.2 grams of chlorine.  Chlorine has a relative mass that is 3.27 times that of element X.  What is the valence of X?

20. How to solve:  First, compare the mass ratios. 39.2/4.0 = 9.8  Next, compare the relative masses. If 9.8 times as much chlorine is needed for every gram of X, divide the 9.8 by the fact that it is 3.27 times heavier,and you get 2.99 which is very close to 3.  This means 3 Chlorines were needed for every X, so X has a valency of 3.

21. Developing “atomic” theory

22. What is an “atom”?  The smallest particle of an element that retains the properties of that element

23. John Dalton (early 1800s) thought atoms were the smallest particles that existed.

24. List Dalton’s postulates  3.When a chemical reaction occurs, atoms c_____, s______ or are re__________.  4. Atoms cannot be subdivided, created or destroyed.  5. When atoms combine they do so in small f____, w____ n____ r____.

25. Are these true?  1. Everything is made up of small particles called atoms  2a.Atoms of different elements _____ in their size, mass, properties.  2a.Atoms of identical elements have _____ size, mass, properties

26. List Dalton’s postulates  3.When a chemical reaction occurs, atoms c_____, s______ or are re__________.  4. Atoms cannot be subdivided, created or destroyed.  5. When atoms combine they do so in small f____, w____ n____ r____.

27. Are these true?  3.When a chemical reaction occurs, atoms c_____, s______ or are re__________.  4. Atoms cannot be subdivided, created or destroyed.  5. When atoms combine they do so in small f____, w____ n____ r____.

28. Now we know that particles smaller than an atom exist.  The first of these particles to be discovered was the “electron”.  J.J. Thomson gets credit for the discovery.

29. What is an “ion”?  A charged particle

30. How was the electron discovered?  Experiments using a CRT Cathode Ray Tube

31. Experiments with the CRT  When an object was placed in the middle, it cast a ______. Meaning:  Something was being blocked.  When a paddle wheel was placed in the middle, it moved to toward the anode. Meaning:  Something was pushing it toward the anode.

32. Experiments with the CRT  When a magnet was near the beam, the beam moved. Meaning:  The stuff in the beam has magnetic properties.  When a negatively charged plate was placed near the tube, the beam moved _____. Meaning:  Whatever is in the beam has a negative charge (because like charges repel.)

33. What is an “ion”?  A charged particle (not neutral)

34. Are all atoms of the same element identical?  No. Different isotopes of each element exist.  Usually one isotope is far more prevalent than others.

35. What is a “isotope”?  A “variety” of an atom that has a specific number of neutrons.  When the same type of atom has a different number of neutrons, that is a different isotope of the same element.  Example C-12, C-13, C-14

36. Are all isotopes stable?  No. Ususally one isotope is stable, and the others are not stable.  Unstable isotopes are “radioactive”.  Example C-12 is a stable isotope,  C-14 is an “unstable” isotope.