1. Foundations of College Chemistry, 14 th Ed. Morris Hein and Susan Arena Lightning occurs when electrons move to neutralize charge difference between the storm clouds and Earth. 5 Early Atomic Theory and Structure Copyright © 2014 John Wiley & Sons, Inc. All rights reserved.

2. 5.1 Dalton’s Model of the Atom 5.2 Electric Charge A. Discovery of Ions 5.3 Subatomic Parts of the Atom 5.4 The Nuclear Atom 5.5 Isotopes of the Elements 5.6 Atomic Mass Chapter Outline © 2014 John Wiley & Sons, Inc. All rights reserved.

3. Early models of the atom were developed by the Greeks. Empedocles proposed matter was composed of four basic elements: earth, air, water and fire. Democritus proposed matter was composed of small, indivisible particles he called atoms. Atoms could combine in different ways, giving rise to the diversity of compounds we observe. Aristotle, an influential philosopher, supported Empedocles’ theory, so atomic theory was not fully accepted until 2000 years later. Early Theories on the Structure of Matter © 2014 John Wiley & Sons, Inc. All rights reserved.

4. Dalton’s theory of atoms, proposed in the early 1800s, states: 1.Elements are composed of small, indivisible particles called atoms. 2.Atoms of the same element are identical in mass and size. 3.Atoms of different elements differ in their mass and size. 4.Compounds are formed by combining two or more atoms of different elements. 5.Atoms combine to form compounds in simple whole number ratios. 6.Atoms of two elements may combine in different ratios, leading to formation of different compounds. Dalton’s Model of the Atom © 2014 John Wiley & Sons, Inc. All rights reserved.

5. a.Atoms are individual particles which are different for each element. b./c. Atoms combine in fixed ratios to form compounds. Two elements can combine in varying ratios to give different compounds. H2O2H2O2 H2OH2O Most of Dalton’s theory remains valid today. Dalton’s Model of the Atom © 2014 John Wiley & Sons, Inc. All rights reserved.

6. 1.Elements can be decomposed under certain conditions. 2.Not all atoms of the same element have identical mass. These are called isotopes. Revisions to Dalton’s Theory 3.Atoms are not indivisible. Atoms are composed of subatomic particles. Dalton’s Model of the Atom © 2014 John Wiley & Sons, Inc. All rights reserved.

7. Properties of Electric Charge 1.Charge may be either positive or negative. 2.Opposite charges (positive and negative) attract while like charges (i.e. negative and negative) repel. 3.Charge may be transferred from one object to another, by contact or induction. 4.The force of attraction between charges (F) is related to the distance between charges by: F = r2r2 kq 1 q 2 where q 1 and q 2 are the charges, r is the distance between charges, and k is a constant. Electric Charge © 2014 John Wiley & Sons, Inc. All rights reserved.

8. Michael Faraday: English scientist who discovered electrolytes (compounds that conduct electricity when dissolved in water). Faraday also discovered that some compounds decompose in water into their elements. These charged elements are called ions. A light bulb glows when ions are present in a saltwater solution when current is passed through it. These elements were attracted to either negatively or positively charged electrodes in the solution, meaning they were no longer neutral. Discovery of Ions © 2014 John Wiley & Sons, Inc. All rights reserved.

9. Arrhenius extended Faraday’s work. He proposed ions are atoms (or groups of atoms) that carry a positive or negative charge. Ex. NaCl in water dissociates into two ions, Na + and Cl –. The Na + (cation) produced is attracted to the negatively charged electrode (cathode). The Cl – (anion) produced is attracted to the positively charged electrode (anode). Based on Faraday’s and Arrhenius’ work, Stoney proposed the electron was a fundamental unit of electricity associated with atoms. J. J. Thomson later experimentally confirmed the existence of electrons. The Nature of Ions © 2014 John Wiley & Sons, Inc. All rights reserved.

10. Because atoms are so small, determining the presence of subatomic particles was very difficult. A single atom is tiny (diameter of 0.1 to 0.5 nm). A scanning tunneling microscope (STM) image shows an array of Cu atoms. New instruments in the early 1900s permitted detection of these particles. Subatomic Parts of the Atom © 2014 John Wiley & Sons, Inc. All rights reserved.

11. A Crooks (cathode) ray tube. The stream of electrons passes between the electrodes. A Crooks tube permits generation of cathode rays, which are streams of electrons. The electron beam is deflected by both electric and magnetic fields, indicating it has charge. Subatomic Parts of the Atom © 2014 John Wiley & Sons, Inc. All rights reserved.

12. Electrons (e – ): A particle with negative electrical charge (assigned a relative charge of –1). Electrons have a very small mass (9.110 x 10 –28 g) and size (<10 –12 cm). Protons (p): A particle with positive electrical charge (assigned a relative charge of +1). Protons have a much larger mass (~1837 times the mass of an electron). Electrons and Protons © 2014 John Wiley & Sons, Inc. All rights reserved.

13. Thomson’s work demonstrated the atom is composed of smaller, charged particles. Dalton’s theory of the atom then had to be revised. Thomson’s Model of the Atom Electrons are negatively charged particles which are embedded in a positively charged atomic sphere. Thomson’s “plum pudding” model of the atom. Electrons + charged sphere The Effect of Subatomic Particles © 2014 John Wiley & Sons, Inc. All rights reserved.

14. Electrons are gained from atoms to give anions. Electrons are lost from atoms to give cations. Atoms can become ions by gaining or losing electrons from this sphere. The Effect of Subatomic Particles © 2014 John Wiley & Sons, Inc. All rights reserved.

15. The last subatomic particle was discovered by Chadwick in 1932. Neutrons (n) A particle with no electrical charge. Neutrons have a mass similar to that of a proton. Neutrons © 2014 John Wiley & Sons, Inc. All rights reserved.

16. Chemical properties of atoms can be described based on the electrons, protons and neutrons. Though other subatomic particles are now known, the theories of atomic structure are based only on these 3 subatomic particles. Atoms are composed of three smaller, subatomic particles: electrons, protons and neutrons. Summary of Subatomic Particles © 2014 John Wiley & Sons, Inc. All rights reserved.

17. In 1911, Ernest Rutherford established the nuclear model of the atom by bombarding gold atoms with α particles. This suggested the gold atoms must have a densely, positively charged nucleus to affect the path of an α particle (a positively charged He atom). Most of the particles passed through the gold foil, but some were deflected and some even bounced back! Nuclear Model of the Atom © 2014 John Wiley & Sons, Inc. All rights reserved.

18. Because most of the particles were not deflected, this suggested most of the atom is empty space. Protons and neutrons are located in the nucleus. Electrons are dispersed throughout the remainder of the atom (mainly open space). Neutral atoms contain the same number of protons and neutrons to maintain charge balance. Nuclear Model of the Atom © 2014 John Wiley & Sons, Inc. All rights reserved.

19. Atomic Number: Number of protons in the nucleus of an atom. Atomic numbers for every element are above the element’s symbol in the periodic table. The atomic number determines the identity of the atom. 27 Co Atomic Number © 2014 John Wiley & Sons, Inc. All rights reserved.

20. After discovery of the nuclear model of the atom, the mass of almost all atoms was found to be larger than expected, based on the number of protons and electrons. This led to the discovery of neutrons. Though all atoms of the same element have the same number of protons, atoms of the same element may have different numbers of neutrons. Isotopes: atoms of an element with the same atomic number but different numbers of neutrons. Isotopes of the Elements © 2014 John Wiley & Sons, Inc. All rights reserved.

21. 1 proton 1 neutron Standard Isotopic Notation Mass number: Total number of protons and neutrons for an element. Example: Isotopes of Hydrogen 1 proton 0 neutrons 1 proton 2 neutrons Protium Deuterium Tritium A E Element Symbol Z Atomic Number Mass Number Isotopes of the Elements © 2014 John Wiley & Sons, Inc. All rights reserved.

22. Practice: How many protons, neutrons, and electrons are found in each of the following isotopes? 64 Cu 29 Atomic Number: 29 protons (therefore 29 electrons) # Neutrons = Mass Number – Atomic Number 64 – 29 = 35 neutrons Isotopes of the Elements © 2014 John Wiley & Sons, Inc. All rights reserved.

23. Which isotope corresponds to an element with 24 protons and 28 neutrons? Solution: 28 Cr 52 Cr 24 52 Ni 28 128 Te 52 a. b. c. d. 24 Cr 52 e. # protons = Atomic Number = 24 Element: Cr Mass Number = protons + neutrons = 24 + 28 = 52 Let’s Practice! © 2014 John Wiley & Sons, Inc. All rights reserved.

24. Because the mass of a single atom is so small, it is inconvenient to use this as a mass unit. Instead, relative atomic mass units (amu) are used. Using carbon-12,, as a standard, 1 atomic mass unit is equal to 1/12th the mass of a carbon-12 atom. 12 C 6 1 amu = 1.6606 x 10 -24 g All periodic tables use atomic masses based on the carbon-12 isotope. Atomic Mass © 2014 John Wiley & Sons, Inc. All rights reserved.

25. Since most elements are a mixture of isotopes, the atomic mass for an element is the weighted average of all naturally occurring isotopes of the element. Example: The atomic mass of Cu is 63.546 amu. Cu exists as 2 major isotopes, Cu-63 and Cu-65. Cu-63 is more abundant, as the atomic mass is very close to 63 amu. Calculating average atomic mass: Sum of the atomic mass of each isotope multiplied by its % abundance. Atomic Mass and Isotope Distribution © 2014 John Wiley & Sons, Inc. All rights reserved.

26. Average atomic mass of Cu: Measuring Cu isotope abundances by using mass spectrometry. (62.9298) x (0.6909) + (64.9278) x (0.3091) = 63.55 amu Atomic Mass % Abundance Atomic Mass and Isotope Distribution © 2014 John Wiley & Sons, Inc. All rights reserved.

27. Silver exists as two isotopes with atomic masses of 106.9041 and 108.9047 amu. Determine the average atomic mass for silver if the % abundance for each isotope is 51.82 and 48.18%, respectively. Average atomic mass of Ag: (106.9041) x (.5182) + (108.9047) x (0.4818) = 107.8680 amu Atomic Mass % Abundance Atomic Mass Practice © 2014 John Wiley & Sons, Inc. All rights reserved.

28. a. 36.95690 d. 34.96885 b.36.57823 e. 33.56438 c. 35.64544 Chlorine exists as two isotopes, Cl-37 (36.96590 amu) and Cl-35. If the percent abundance of each isotope is 24.47 % and 75.53 %, what is the atomic mass of Cl-35 if the average atomic mass is 35.46 amu? (36.96590) x (.2447) + (a) x (0.7553) = 35.46 amu Solve for a: Solution: 9.046 + (a) x (0.7553) = 35.46 amu (a) x (0.7553) = 26.41 amu a = 34.97 amu Let’s Practice! © 2014 John Wiley & Sons, Inc. All rights reserved.

29. Describe Dalton’s model of the atom and compare it to the earlier concepts of matter. 5.1 Dalton’s Model of the Atom Use Coulomb’s Law to calculate the force between particles and distinguish between a cation and anion. 5.2 Electric Charge Describe the three basic subatomic particles and how they changed Dalton’s model of the atom. 5.3 Subatomic Parts of the Atom Learning Objectives © 2014 John Wiley & Sons, Inc. All rights reserved.

30. Explain how the nuclear model of the atom differs from the Dalton and Thomson models. 5.4 The Nuclear Atom Define the terms atomic number, mass number and isotope. 5.5 Isotopes of the Elements Define the relationship between the atomic mass of an element and the masses of its isotopes. 5.6 Atomic Mass Learning Objectives © 2014 John Wiley & Sons, Inc. All rights reserved.