Atoms and Elements

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

2. The Atomic Theory of Matter • 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. (e.g., H2O, C6H6, C12H22O11 but not H2, Cl2)

3. Law of Constant Composition In a compound, relative amounts and kinds of atoms are fixed. For example, water (H2O) is always a ratio of two hydrogen atoms to one oxygen atom. 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 H2O).

4. The Discovery of Atomic Structure 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. A charged particle will have its path bend in either an electric or magnetic field. 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.

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

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

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

8. The Discovery of Atomic Structure 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.

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

10. The Discovery of Atomic Structure 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. 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. The Discovery of Atomic Structure 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.

12. The Modern View of Atomic Structure The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). • 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. The Nucleus Ångstrom unit: (1Å = 10-8cm =10-10 m)

14. The Modern View of Atomic Structure Isotopes, Atomic Numbers, and Mass Numbers Atomic number (Z) = number of protons in the nucleus. Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). By convention, for element X, we write Isotopes have the same Z but different A. Isotopes of carbon: 116C, 126C, 136C, 146C Note the common “6” for all isotopes of carbon

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. in the nucleus about the same 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

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

18. Exercise

19. Exercise

20. Atomic and Molecular Weights The Atomic Mass Scale Assume H has a relative mass of 1 “unit” H2O 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. Another example: 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 Then, mass C = 12 (Show this!) mass H 1

22. Mass of 12C = 12 amu exactly 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 12C having mass of 12 exactly (by definition). This unit of mass is called the atomic mass unit, or amu.

23. Atomic and Molecular Weights Average Atomic Mass Because of isotopes, atoms have average atomic weights Relative atomic mass: average masses of isotopes: Naturally occurring C: 98.892 %12C + 1.108 %13C. 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.

24. Formula and Molecular Weights 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(C6H12O6) = 6(12.0 amu) + 12(1.0 amu) + 6(16.0 amu) = 180.0 amu We will use FW and MW interchangeably

25. The Periodic Table

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

27. The Periodic Table Columns in the periodic table are called groups (numbered from 1A to 8A or 1 to 18). Rows in the periodic table are called periods. 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.

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

29. Some elements occur naturally as diatomicmolecules (Most elements can be viewed as uniatomic; but there are unusual elemental molecules, e.g., P4, S8, C60.)

30. The Periodic Table Some of the groups in the periodic table are given special names. • 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. Halogens Chalcogens Alkaline Earths Alkali Metals Navigating the Periodic Table Transition Metals Noble or Inert Gases Lanthanides (rare earths) Actinides

32. Different Kinds of Compounds A salt, formed by ionic bonding, is formed between a metal and a nonmetal, (e.g., NaCl, Ag2O).

33. Different Kinds of Compounds A molecule, formed by covalent bonding, is formed between a nonmetal and a nonmetal, (e.g., CO2, PBr3, H2O).

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. The Mole The mole connects the visible with the invisible. A fluorine molecule (F2) 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 1023. We will be using the mole concept very often. Amedeo Avogadro 1776-1856

36. The Mole Examples: A mole of H is 1.008 grams. A mole of H2 is 2.016 grams. A mole of CO2 is 44.011 grams. A mole of CO is 28.01 grams. A mole of octane (C8H18) 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 (NaHCO3) is 84.01 grams. A mole of Ag2O is 231.74 grams. A mole of glucose (C6H12O6) is 180.16 grams. A mole of chlorophyll (C55H72MgN4O5) is 893.51 grams.