General Chemistry I Virginia State University

1. General Chemistry IVirginia State University Chapter 2 Dr. Vilchiz Fall 2013

2. THE UNIVERSE • According to Science how did it all started? • Evidence? • Direct? • Indirect?

3. Questions… • How big can it get? • How small was it? • How much mass is there? • Etc.

4. Measurements • Types of Properties • Types of Measurements • Parts in a measurement • Units • Basic vs derived • SI vs Imperial

5. SI UNITS • The SI (Systeme Internationale) of units has six basic units • Length meter m • Temperature Kelvin K • Mass kilogram kg • Time second s • Luminous Intensity candela cd • Substance amount Mole mol

6. Derived Units • The rest of the units are derived from these basic units. • Volume m3 • Force Newton N=kgm/s2 • Energy Joule J=kgm2/s2 • Pressure Pascal Pa=kg/ms2 • Velocity m/s • Density kg/m3

7. Metric System Similar to Table 2.4 In your book

8. Temperature • What is temperature? • What different scales do we use? • Where did they come from?

9. Scientific Notation • What is it? • Why use it? • What else does it accomplish? • What practical benefits does it have?

10. Measurement and Sig Figs • All measurements, no matter how careful they are taken, have an error, uncertainty, associated with them.

11. Errors • Type of errors? • How can we “fixed” them? • How do they affect reported numbers? • Accuracy vs Precision

12. Rules for Sig Figs • What numbers count? • All nonzero digits are significant. • Zeros between significant figures are significant. • Zeros preceding the first nonzero digit are not significant. • Zeros to the right of the decimal after a nonzero digit are significant. • Zeros at the end of a nondecimal number may or may not be significant. (Use scientific notation.)

13. Significant Figures 3456has 4sig figs. 0.0486 has 3 sig figs. 16.07has 4sig figs. 1.300has 4sig figs. 1310has 3sig figs. 1310.has 4sig figs.

14. Significant Figures & Math • When multiplying and dividing measured quantities, give as many significant figures as the least found in the measurements used. • When adding or subtracting measured quantities, give the same number of decimals as the least found in the measurements used.

15. 14.0 g /102.4 mL = 0.136718g/mL Significant Figures

16. 14.0 g /102.4 mL = 0.136718g/mL Significant Figures only three significant figures

17. Significant Figures 14.0 g /102.4 mL = 0.137 g/mL only three significant figures

18. Exact Numbers • What are they? • Do they have sig figs? • How do they affect calculations?

19. Calculations… • Dimensional Analysis • Conversion Factors • How accurate do I have to be?

20. The Electron • The electron was discovered by J.J. Thomson between 1898 and 1903. NP 1906 • While an electrical discharge is applied to a tube filled with gas he noticed a “ray” (cathode ray tube). If a magnetic field was present near the ray it bent the ray away from the negative pole. • The ray was made of negative particles (electrons).

21. A cathode-ray tube. The fast-moving electrons excite the gas in the tube, causing a glow between the electrodes.

22. Deflection of cathode rays by an applied electric field.

23. The plum pudding model of the atom.

24. Charge of an Electron • We all know that the electron is negatively charged… but just how much charged does it have? • Robert Millikan devised an experiment that answer this question.

25. Millikan’s Experiment

26. The Nucleus • In 1911 Rutherford performed an experiment in which he bombarded a thin piece of gold foil with alpha particles. • An alpha particle is a helium nuclei. • He noticed that while most of the alpha particles went through some were reflected/deflected. The only explanation to this was that there was a solid component in the inside of the gold atoms.

27. Rutherford's experiment on -particle bombardment of metal foil.

28. (a) The expected results of the metal foil experiment if Thomson's model were correct. (b)Actual results.

29. Composition of an Atom • The nucleus of an atom is composed of two different kinds of particles: protons and neutrons. • Since they reside in the nucleus sometimes they are referred as nucleons. • The nucleus is surrounded by electron clouds which are much bigger than the nucleus. • Electron clouds have no definite shape and therefore are hard to measure in size.

30. Atomic Theory of Matter • A proton is the nuclear particle having a charge equal in magnitude to that of the electron’s (a “unit” charge) but opposite in sign. The proton’s mass is more than 1800 times that of the electron’s yet it is 10000 times smaller. (representation of an atom) • The number of protons in the nucleus of an atom is referred to as its atomic number (Z).

31. Atomic Theory of Matter • An element is a substance whose atoms all have the same atomic number. This means that the number of protons dictates the type of element. • The neutron is a nuclear particle having a mass almost identical to that of a proton, but no electric charge. Notice there is no restriction as to the number of neutrons for a given element. • Table 2.1 summarizes the masses and charges of these three fundamental particles.

32. Atomic Theory of Matter • The mass number is the total number of protons and neutrons in a nucleus. • A nuclide is an atom characterized by a definite atomic number and mass number. • The shorthand notation for a nuclide consists of its symbol with the atomic number as a subscript on the left and its mass number as a superscript on the left.

33. Nuclear Symbols As we can see an atom belonging to a given element may have different compositions. So we must figure out how to represent the different possible configurations.

34. Nuclear Symbols In a nuclear symbol the Z represents the atomic number (# of protons which then dictates X) The X represents the element’s symbol (periodic table) The Y represents the charge of the species (difference between p+s and e-s) The A represents the atomic mass (sum of p+s and ns)

35. Isotopes Atoms that belong to the same element but have different compositions are called isotopes. In the case of Hydrogen ~98% of the hydrogen is present as 1H, ~1.5% as 2H, and <0.5% as 3H. All 3 have 1 proton, 1 electron but the number of neutrons varies.

36. Atomic Theory of Matter • Isotopes are atoms whose nuclei have the same atomic number but different mass numbers; that is, the nuclei have the same number of protons but different numbers of neutrons. • Sodium, for example, exists as two isotopes: sodium-23 and sodium-24. • The fractional abundance is the fraction of a sample of atoms that is composed of a particular isotope.

37. Atomic Weights • The atomic weight of an element is obtained by adding the products of each isotope’s fractional abundance and its isotopic mass. • Example: Calculate the atomic weight of boron, B, from the following data: • ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE • B-10 10.013 0.1978 • B-11 11.009 0.8022

38. Atomic Weights • Calculate the atomic weight of boron, B, from the following data: • ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE • B-10 10.013 0.1978 • B-11 11.009 0.8022 • B-10: 10.013 x 0.1978 = 1.9805 • B-11: 11.009 x 0.8022 = 8.8314 • 10.8119 = • 10.812 amu • ( = atomic wt.)

39. Atomic Weights • Dalton chose Hydrogen to have a mass of 1 “Dalton, ”since he believed it to be the lightest element. • He found that carbon weighed 12 times more than hydrogen. He therefore assigned carbon a mass of 12 “Daltons.”

40. Atomic Weights • Dalton’s atomic weight scale was eventually replaced in 1961, by the present carbon–12 mass scale. • One atomic mass unit (amu) is, therefore, a mass unit equal to exactly 1/12 the mass of a carbon–12 atom. • On this modern scale, the atomic weight of an element is the weighted average atomic mass for the naturally occurring element, expressed in atomic mass units.

41. Atomic Weights • Calculate the atomic weight of boron, B, from the following data: • ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE • B-10 10.013 0.1978 • B-11 11.009 0.8022 • B-10: 10.013 x 0.1978 = 1.9805 • B-11: 11.009 x 0.8022 = 8.8314 • 10.8119 = • 10.812 amu • ( = atomic wt.)

42. Atomic Weights • Dalton chose Hydrogen to have a mass of 1 “Dalton, ”since he believed it to be the lightest element. • He found that carbon weighed 12 times more than hydrogen. He therefore assigned carbon a mass of 12 “Daltons.”

43. Atomic Weights • Dalton’s atomic weight scale was eventually replaced in 1961, by the present carbon–12 mass scale. • One atomic mass unit (amu) is, therefore, a mass unit equal to exactly 1/12 the mass of a carbon–12 atom. • On this modern scale, the atomic weight of an element is the weighted average atomic mass for the naturally occurring element, expressed in atomic mass units.

44. Isotope Composition

45. GENERAL CHEM I The Periodic Table Fall 2013 Dr. Victor Vilchiz

46. The Periodic Table • How is the Periodic Table constructed? • There are many ways to answer this question. • But the real answer might be the most obvious. • PERIODICALLY!!! • Ok so… Periodically but what does it mean to be periodical?

47. What is a period? • A period is something that repeats itself in a given interval. • We will talk more about periods in the next section. • The periodic table can be said to be organized by the number of protons in the nucleus of elements. (Atomic Number) • It can also be said that it is arranged more or less by atomic mass.

48. What is the real answer? • As a Physical Chemist in the 21st Century I can tell you that it is arranged according to the electron configuration of the elements. • Uhm can you pass that through me one more time? • The periodic table is arranged according to the number of electrons in the outermost shell in an atom of each element. • For the average person that means what? • Ok, Ok… it has to do with the number of electrons.

49. That’s Yiddish to me!!! • Imagine if I have my original response back in the 1800’s!!! • I would had been handed my Hemlock and told to make a toast to Socrates. • There were no electrons back then • Lets go back and take it from the 1800’s forward. • Periodic = there are patterns • Lets take several elements and react them.

50. Original Periodic Chart • Take for example Sodium and react it with any other element you can find. • Uhm… Na and a series of other elements ( F, Cl, Br, I) react in a 1:1 ratio. • Therefore, F, Cl, I and Br must be grouped. • Replace Na with K, Li or Cu and the same is true. • Therefore, Na, Li, Cu and K belong in the same group.