1. 1 Atoms and Elements Chapter 2 John Dalton 1766-1844. Conceived atomic weights. Ernest Rutherford 1871-1937.* Discoverer of atomic nucleus. D. I. Mendeleev 1834-1907. Periodic Table. Predicted elements.

2. 2 John Dalton: Elements are composed of atoms. –All atoms of an element are identical (chemically). (Dalton stressed “identical in weight” but he didn’t know about isotopes) –In chemical reactions, the atoms are not changed. –Compounds are formed when atoms of more than one element combine. The Atomic Theory of Matter (e.g., H 2 O, C 6 H 6, C 12 H 22 O 11 but not H 2, Cl 2 )

3. 3 For example, water (H 2 O) is always a ratio of two hydrogen atoms to one oxygen atom. Law of Constant Composition In a compound, relative amounts and kinds of atoms are fixed. Specifically, water is always 11.1% hydrogen and 88.9% oxygen (by weight ). For water, O/H = 88.9/11.1 = 8/1 (by weight) Dalton perceived that this constant ratio of mass of different elements in a compound reflected a specific ratio of atoms, with each element having its specific, but unique weight. But he was confused about the exact number of atoms in a compound. For example, he thought water had a formula of HO (and not H 2 O).

4. 4 The ancient Greeks were the first to postulate that matter consists of indivisible constituents. Later scientists realized that the atom consisted of charged (+ or -) entities. The Discovery of Atomic Structure Cathode Rays and Electrons A cathode ray tube (CRT) is a hollow vessel with an electrode at either end. A high voltage is applied across the electrodes. A charged particle will have its path bend in either an electric or magnetic field.

5. 5 Cathode Rays and Electrons The Discovery of Atomic Structure (electrons are charged (-) particles)

6. 6 The Discovery of Atomic Structure a spot which is not affected by the electric field, Three spots are noted on the detector: a spot in the direction of the positive (+) plate, a spot in the direction of the negative (-) plate.

7. 7  -radiation: Large deflection toward the positive plate corresponding to radiation which is negatively charged and of low mass. These  particles are light and of low mass.  particles are electrons.  -radiation: No deflection; neutral (zero charge) radiation.  -radiation: Small deflection toward the negative plate corresponding to high mass, positively charged radiation.

8. 8 The Nuclear Atom From the separation of radiation we conclude that the atom consists of neutral, positively, and negatively charged entities. J. J. Thomson assumed all these charged species were found in a sphere. The Discovery of Atomic Structure

9. 9 The Nuclear Atom Rutherford’s  -particle experiment: The Discovery of Atomic Structure

10. 10 The Nuclear Atom In order to get the majority of  -particles through a piece of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge  the electron. The Discovery of Atomic Structure To account for the small number of high deflections of the  -particles, the center or nucleus of the atom must consist of a dense positive charge.

11. 11 The Nuclear Atom Rutherford modified Thomson’s model as follows: assume the atom is spherical but a massive positive charge must be located at the center, with a diffuse light negative charge surrounding it. The Discovery of Atomic Structure

12. 12 The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). The Modern View of Atomic Structure Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus. Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.

13. 13 (1Å = 10 -8 cm =10 -10 m) The Nucleus Ångstrom unit:

14. 14 Atomic number (Z) = number of protons in the nucleus. The Modern View of Atomic Structure Isotopes, Atomic Numbers, and Mass Numbers By convention, for element X, we write Isotopes have the same Z but different A. Isotopes of carbon: 11 6 C, 12 6 C, 13 6 C, 14 6 C Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). Note the common “6” for all isotopes of carbon

15. 15 John Dalton 1766-1844. Henry Moseley 1887-1915. Frederick Soddy 1877-1956. Dalton conceived atomic weights 200 years ago. Moseley defined atomic number, using X-ray diffraction, 100 years later. Soddy defined isotopes shortly after.

16. 16 Comparison of Proton, Neutron and Electron Relative ParticleChargeMass (amu) Proton 1+ 1.0073 Neutron neutral 1.0087 Electron 1- 5.486 x 10 -4 1 amu = 1.66 x 10 -24 g about the same in the nucleus

17. 17 Isotopes, Atomic Numbers, and Mass Numbers The Modern View of Atomic Structure

18. 18 Exercise

19. 19 Exercise

20. 20 The Atomic Mass Scale Atomic and Molecular Weights Assume H has a relative mass of 1 “unit” H 2 O is 88.9 % O and 11.1% H (by mass) Mass ratio of O to H (in water) is: 88.9/11.1 = 8/1. Since there are two H for each O, mass ratio of O to H must be 16/1 or: mass O = 16 mass H 1

21. 21 Another example: Then, mass C = 12 (Show this!) mass H 1 CO is 42.9% C and 57.1% O (by mass) Mass ratio of O to C= 57.1/42.9 = 4/3 or: mass O = 4 mass C 3 Since: mass O = 16 (previous slide) mass H 1

22. 22 Atomic and Molecular Weights Can now build up “relative” atomic mass scale. Roughly, if H=1, then C=12 and O=16. Can add other elements (e.g., N=14, F=19, etc). Current scale is based on isotope 12 C having mass of 12 exactly (by definition). This unit of mass is called the atomic mass unit, or amu. Mass of 12 C = 12 amu exactly

23. 23 Because of isotopes, atoms have average atomic weights Relative atomic mass: average masses of isotopes: Naturally occurring C: 98.892 % 12 C + 1.108 % 13 C. Average mass of C: (0.98892)(12 amu) + (0.0108)(13.00335) = 12.011 amu. Atomic weight (AW) is also known as average atomic mass (atomic mass). These average atomic weights, as found in the earth’s crust, are listed on the periodic table. Atomic and Molecular Weights Average Atomic Mass

24. 24 Formula weights (FW): sum of AW for atoms in formula. FW (NaCl) = AW(Na) + AW(Cl) = 23.0 amu+ 35.5 amu= 58.5 amu (Remember, NaCl is not a “molecule” It exists as a 3-D array of ions) Molecular weight (MW) is the weight of the molecular formula. MW(C 6 H 12 O 6 ) = 6(12.0 amu) + 12(1.0 amu) + 6(16.0 amu) = 180.0 amu Formula and Molecular Weights We will use FW and MW interchangeably

25. 25 The Periodic Table

26. 26 The Periodic Table is used to organize the 114 elements in a meaningful way The Periodic Table As a consequence of this organization, there are periodic properties associated with the periodic table.

27. 27 Columns in the periodic table are called groups (numbered from 1A to 8A or 1 to 18). The Periodic Table Metals are located on the left hand side of the periodic table (most of the elements are metals). Non-metals are located in the top right hand side of the periodic table. Elements with properties similar to both metals and non-metals are called metalloids and are located at the interface between the metals and non-metals. Rows in the periodic table are called periods.

28. 28 The Periodic Table Metals Non-Metals Metalloids “semiconductors”

29. 29 Some elements occur naturally as diatomic molecules (Most elements can be viewed as uniatomic; but there are unusual elemental molecules, e.g., P 4, S 8, C 60.)

30. 30 Some of the groups in the periodic table are given special names. The Periodic Table These names indicate the similarities between group members: Group 1A: Alkali metals - “al kali” = “the ashes” (of a fire) Group 2A: Alkaline earth metals (“earths” historically were oxides that were difficult to reduce to the metal). Group 6A: Chalcogens - “ore formers” Group 7A: Halogens - “salt formers” Group 8A: Noble gases - “unreactive” gases At the bottom are the lanthanides (“rare earths”) and the actinides.

31. 31 Alkali Metals Alkaline Earths Noble or Inert Gases HalogensChalcogens Lanthanides (rare earths) Actinides Transition Metals Navigating the Periodic Table

32. 32 Different Kinds of Compounds A salt, formed by ionic bonding, is formed between a metal and a nonmetal, (e.g., NaCl, Ag 2 O).

33. 33 A molecule, formed by covalent bonding, is formed between a nonmetal and a nonmetal, (e.g., CO 2, PBr 3, H 2 O). Different Kinds of Compounds

34. 34 Different Kinds of Compounds An alloy, formed by metallic bonding, is formed between a metal and a metal, (e.g., brass or nickel-steel)

35. 35 The Mole The mole connects the visible with the invisible. A fluorine molecule (F 2 ) weighs 38.000 amu. A mole of fluorine molecules weighs 38.000 grams. The number of fluorine molecules in a mole is an incredibly large number, called Avogadro’s Number, N, which is 6.022 x 10 23. We will be using the mole concept very often. Amedeo Avogadro 1776-1856

36. 36 The Mole Examples: A mole of H is 1.008 grams. A mole of H 2 is 2.016 grams. A mole of CO 2 is 44.011 grams. A mole of CO is 28.01 grams. A mole of octane (C 8 H 18 ) is 114.22 grams. A mole of copper (Cu) is 63.54 grams. A mole of table salt (NaCl) is 58.44 grams. A mole of sodium bicarbonate (NaHCO 3 ) is 84.01 grams. A mole of Ag 2 O is 231.74 grams. A mole of glucose (C 6 H 12 O 6 ) is 180.16 grams. A mole of chlorophyll (C 55 H 72 MgN 4 O 5 ) is 893.51 grams.