Foundations of Atomic Theory and the Structure of the Atom

1. Unit 3 Atoms, Sub-Atomic Particles & Nuclear Chemistry

2. The Particle Theory of Matter • In 400 B.C. Democritus, a Greek philosopher, first proposed the idea of a basic particle of matter that could not be divided any further. • He called this particle the atom, based on the Greek word atomosmeaning indivisible. • This early theory was not backed up by experimental evidence and was ignored by the scientific community for nearly 2000 years.

3. Foundations of Atomic Theory • By the late 1700s, experiments with chemical reactions led to the discovery of 3 basic laws: • The Law of Conservation of Mass. • This law states that mass is neither created nor destroyedduring ordinarychemical reactions or physical changes. • Formulated by Antoine Lavoisier in 1789.

4. Foundations of Atomic Theory The Law of Conservation of Mass

5. Foundations of Atomic Theory • The Law of Definite Proportions • States that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. • Discovered by Joseph Proust in 1797. Visual Concept

6. Foundations of Atomic Theory • The Law of Multiple Proportions – • States that if different compounds are composed of the same 2 elements, then the ratio of the masses of the elements is always a ratio of small whole numbers. • Published by John Dalton in 1804.

7. Dalton’s Atomic Theory • In 1808, John Dalton proposed an explanation for the three laws. His atomic theory states: • All matter is composed of atoms. • Atoms of the same element are identical; atoms of different elements are different. • Atoms cannot be subdivided, created, or destroyed. • Atoms of different elements combine in simple whole-number ratios to form chemical compounds. • In chemical reactions, atoms are combined, separated, or rearranged.

8. Corrections to Dalton’s Theory • Dalton turned Democritus’s idea into a scientific theory that could be tested by experiment. • But not all aspects of Dalton’s theory have proven to be correct. We now know that: • Atoms are divisible into even smaller particles. • A given element can have atoms with different masses.

9. Dalton’s Atomic Model • An atom is the smallest particle of an element that has all the properties of that element. • Atoms are too small to see…even through the most powerfulmicroscope!! • Dalton thought atoms were solid balls of matter and were indivisible.

10. Discovery of the Electron • In 1897, Joseph John (JJ) Thomson showed that cathode rays are composed of identical negatively charged particles, which were named electrons. • The electron was the first subatomic particle to be discovered.

11. Thomson’s Cathode Ray Tube Experiment Visual Concept

12. Charge and Mass of the Electron • In 1909, Robert Millikan measuredthe charge on the electron during his oil drop experiment. • Using the charge-to-mass ratio, scientists were able to figure out the mass of the electron: about 1/2000 the mass of a hydrogen atom.

13. Millikan’s Oil Drop Experiment Visual Concept

14. Thomson’s Plum Pudding Model • After the work of Thomson andMillikan, the accepted model of the atom was called the plum pudding model. • The atom was viewed as a ball of positively- charged material with tiny negatively-charged electrons spread evenly throughout.

15. Discovery of the Atomic Nucleus • More detail of the atom’s structure was provided in 1911 by Ernest Rutherford and his associates Hans Geiger and Ernest Marsden. • The results of their gold foil experiment led to the discovery of a very densely packed bundle of matter with a positive electric charge. • Rutherford called this positive bundle of matter the nucleus.

16. The Gold Foil Experiment Visual Concept

17. Rutherford’s Atomic Model • After Rutherford’s gold foil experiment, the accepted model of the atom looked like this: • A small, positively-charged nucleus with negative electrons surrounding it at some distance away. Most of the atom is empty space.

18. Structure of the Atom • Rutherford proposed that the nucleus had particles with the same amount of charge as an electron but the opposite sign, called protons. • Relative charge = +1 • Relative mass = 1 amu • For an atom to be neutral there must be equal numbers of protons and electrons. • Throughout the 1920’s scientists accepted an (incorrect) model of the atom composed of protons and electrons.

19. Some Problems • How could beryllium have 4 protons stuck together in the nucleus? • shouldn’t they repel each other? • If a beryllium atom has 4 protons, then it should weigh 4 amu; but it actually weighs 9.01 amu! Where is the extra mass coming from? • Each proton weighs approximately 1 amu. • The electron’s mass is only about 0.00055 amu and Be has only 4 electrons, so they don’t account for the extra 5 amu of mass.

20. There Must Be Something Else There! • These questions were answered in 1932 by James Chadwick (a student of Rutherford’s), who discovered another particle in the nucleus, which he called a neutron. • Charge = 0 (no charge). • Relative mass = 1 amu.

21. Subatomic Particles • The nucleus is made up of at least one positively charged particle called a protonand usually one or more neutral particles called neutrons. • Protons, neutrons, and electrons are often referred to as subatomic particles.

22. Atomic Number • The atomic number (Z) of an element is the number of protons of each atom of that element. • Atoms of the same element all have the same number of protons.

23. Mass Number • The mass numberis the total number of protons and neutrons in the nucleus of an atom. • Atoms of the same element can havedifferentmass numbers.

24. Isotopes • Isotopes are atoms of the same element that have different masses. • Isotopes have the same number of protons and electrons but different numbers of neutrons. • Most of theelements consist of mixtures ofisotopes.

25. Designating Isotopes • Hyphen notation: The mass number is written with a hyphen after the name of the element. uranium-235 • Nuclear symbol:The superscript indicates the mass number and the subscript indicates the atomic number. Mass number Atomic number

26. Calculating Neutrons • The number of neutrons is found by subtracting the atomic number from the mass number. mass number − atomic number = number of neutrons • Nuclideis a general term for a specific isotope of an element.

27. Calculating Subatomic ParticlesSample Problem How many protons, electrons, and neutrons are there in an atom of chlorine-37? Solution: Number of protons Number of electrons Number of neutrons = atomic number (on periodic table) 17 = number of protons 17 = mass number - protons 20

28. The Atomic Mass Unit • The standard used by scientists to compare units of atomic mass is the carbon-12 atom. • One atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom. • The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.

29. Average Atomic Mass • Average atomic massis the weighted average of the atomic masses of the naturally occurring isotopes of an element. • The average atomic mass of an element depends on both the massand the relative abundance of each of the element’s isotopes.

30. Calculating a Weighted Average If you have the following grades, what would your marking period average be? • First, change percents to decimals. • Next, multiply each grade by its decimal percent. • Finally, add up all the products. ÷ 100 = 0.45 x = 32.4 x = 7.8 ÷ 100 = 0.10 x = 21.0 ÷ 100 = 0.25 x = 9.8 ÷ 100 = 0.10 x = 9.4 ÷ 100 = 0.10 + 80.4

31. Calculating Average Atomic MassSample Problem 1 Copper consists of 69.15% copper-63, which has an atomic mass of 62.929 601 amu, and 30.85% copper-65, which has an atomic mass of 64.927 794 amu. What is the Average Atomic Mass of Copper? Solution • Change percents to decimals. • Multiply the atomic mass of each isotope by its relative abundance. • Add up all of the products. x = 43.52 x = 20.03 + 63.55

32. Calculating Average Atomic MassSample Problem 2 A student believed that she had discovered a new element and named it mythium. Analysis found it contained two isotopes. The composition of the isotopes was 19.9% of atomic mass 10.013 and 80.1% of atomic mass 11.009. What is the average atomic mass, and do you think mythium was a new element? Solution: • Average Atomic Mass:(.199 x 10.013) + (.801 x 11.009) = 10.811 • Because the atomic mass is the same as the atomic mass of boron, mythium was not a new element. Round off to: 10.8

33. Forces in the Atom • Electrons and protons attract because of opposite electrical charges, but protons and protons repel since they have the same charge. • The nucleus is heldtogether by a mysterious force called the strong nuclear force which only exists between nucleons (protons and neutrons) whichare very close together.

34. Valley of Stability for Z = 1 20, stable N/Z ≈ 1 for Z = 20  40, stable N/Z approaches 1.25 for Z = 40  80, stable N/Z approaches 1.5 for Z > 83, there are no stable nuclei

35. Naturally Radioactive Elements • All of the elements beyond atomic number 83 are unstable and thus radioactive.

36. Nuclear Reactions • Large, unstable nuclei spontaneously break apart to form smaller, more stable nuclei. • A nuclear reaction is a reaction that affects the nucleus of an atom.Example: • A transmutation is a change in the identity of a nucleus as a result of a change in the number of its protons.

37. Nuclear ReactionsSample Problem Identify the products that balance the following nuclear reactions: a. b. Solution: • Atomic Mass: 212 = 4 + _____ Atomic Number: 84 = 2 + _____ • Atomic Mass: 22 + ____ = 22 Atomic Number: 11 + ____ = 10 208 82 0 -1

38. Radioactive Decay • Radioactive decay is the spontaneous disintegration of a nucleus into a lighter nucleus, accompanied by nuclear radiation. • Nuclear radiation is particles and/or electromagnetic radiation emitted from the nucleus during radioactive decay.

39. Types of Radioactive Decay Alpha Emission • An alpha particle (α) is two protons and two neutrons bound together and is emitted from the nucleus during some kinds of radioactive decay. • The atomic number decreases by two and the mass number decreases by 4.

40. Types of Radioactive Decay (continued) Beta Emission • Abeta particle (β) is an electron emitted from the nucleus during some kinds of radioactive decay (a neutron can be converted into a proton and an electron.) • The atomic number increases by one and the mass number stays the same.

41. Types of Radioactive Decay (continued) Positron Emission • A positron (β+) is a particle that has the same mass as an electron, but has a positive charge (to decrease the number of protons, a proton can be converted into a neutron by emitting a positron.) • The atomic number decreases by one and the mass number stays the same.

42. Types of Radioactive Decay (continued) Electron Capture • In electron capture, an inner orbital electron is captured by the nucleus of its own atom. (An inner orbital electron combines with a proton to form a neutron.) • The atomic number decreases by one and the mass number stays the same.

43. Types of Radioactive Decay (continued) Gamma Emission • Gamma rays () are high-energy electromagnetic waves emitted from an unstable nucleus. • Atomic number and mass number both stay the same because gamma rays have no charge and no mass. • They are pure energy, and very dangerous to living things.

44. Comparing Alpha, Beta and Gamma • Alpha particles are big andslow. They can’t penetrate skin or paper. • Beta particles have about 100 x the penetrating power of alpha. They can be stopped by clothing, wood, or aluminum foil. • Gamma rays have the greatest penetrating ability. They can only be stopped by a thick layer of lead or concrete. They cause a lot of damage to living cells. Visual Concept

45. Half-Life • Half-life is the time required for half the atoms of a radioactive nuclide to decay. • Each radioactive nuclide has its own half-life. • More-stable nuclides decay slowly and have longer half-lives.

46. Half-LifeSample Problem 1 The half-life of radon-222 is 4 days. After what time will ¼ of a given amount of radon remain? Solution: Determine the number of half-lives that it takes to cut a sample to ¼ of the original amount. Then multiply that number by the half-life. 8 days 2 x 4 days =

47. Half-LifeSample Problem 2 Phosphorus-32 has a half-life of 14.3 days. How many milligrams of phosphorus-32 remain after 57.2 days if you start with 4.0 mg of the isotope? Solution: Determine the number of half-lives in 57.2 days. For each half-life, multiply the original amount by ½: 57.2 days 4 half-lives ÷ 14.3 days = x ½ x ½ 0.25 mg 4.0 mg x ½ = x ½

48. Uses of Radiation • Radioactive dating – scientists can determine the approximate age of an object based on the amount of certain radioactive nuclides present.

49. Uses of Radiation (continued) • Radioactive tracers are radioactive atoms that are incorporated into substances so that movement of the substances can be followed by radiation detectors. • Radioactive tracers can be used by doctors to diagnose diseases. • Radioactive tracers are also used in agriculture to determine the effectiveness of fertilizers.

50. Uses of Radiation (continued) • Irradiated Food – nuclear radiation is used to prolong the shelf life of food.