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Chapter 4

Chapter 4. Atomic Structure. Before the atom…. Many cultures believed that all things were composed of the classical elements: Earth, Wind, Water, and Fire Much of their understanding came from things that they could see – they could not see atoms. 4.1 Atoms.

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Chapter 4

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  1. Chapter 4 Atomic Structure

  2. Before the atom…. • Many cultures believed that all things were composed of the classical elements: • Earth, Wind, Water, and Fire • Much of their understanding came from things that they could see – they could not see atoms.

  3. 4.1 Atoms • The discovery and study of atoms (the fundamental unit of all matter) has been done through indirect observations, logic, and scientific deduction. • These particles can only be viewed with the most powerful telescopes.

  4. 4.1 Atoms • Atoms were first proposed to exist by a Greek philosopher, Democritus, more than 2000 years ago. • Democritus thought that atoms were invisible, indestructible, fundamental units of matter. • His ideas agreed with later scientific theories, but lacked experimental support because scientific experiments were unknown in his time.

  5. 4.1 Atoms • 2200 years after Democritus, John Dalton (1766-1844), proposed an atomic theory. • Dalton, who studied chemistry very differently than Democritus, performed experiments to arrive at his atomic theory.

  6. 4.1 Daltons Atomic Theory 1. All elements are composed of submicroscopic, indivisible particles called atoms. 2. Atoms of the same element are identical. • The atoms of any one element are different from those of any other element.

  7. 4.1 Dalton’s Atomic Theory 3. Atoms of different elements can physically mix together or can chemically combine with one another in simple whole-number ratios to form compounds. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. • However, atoms of one element are never changed into atoms of another element as a result of a chemical reaction.

  8. 4.1 Atoms • atom – the smallest particle of an element that retains the properties of that element. • Example: Pure copper “penny” would contain approximately 2.4 x 1022 copper atoms – that is 3.5 trillion times the number of people on this planet.

  9. 4.1 Atoms • Even though atoms are so small, we can see images from them with the right technology. • Scanning Tunneling Microscopes (STMs) and Atomic Force Microscopes (AFMs) can be used to visualize individual atoms. TEAM Microscope TEAM Published Article

  10. 4.1 Concept Practice 1. Democritus and Dalton both proposed that matter consists of atoms. Explain how their approaches to reaching the same conclusion differed?

  11. 4.1 Concept Practice 2. Which of these statements would Dalton have agreed with (use Dalton’s atomic theory)? a. Atoms are the smallest particles of matter b. The mass of an iron atom is different from the mass of a copper atom c. Every atom of silver is identical to every other atom of silver d. A compound is composed of atoms of two or more different elements

  12. 4.2 Electrons, Protons, Neutrons • Most of Dalton’s Atomic Theory is accepted today – one important revision is that atoms are not indivisible (they can be broken down) • Protons, neutrons, and electrons are three subatomic particles that make up an atom. • There are dozens of subatomic particles, but we will only be studying these three subatomic particles in chemistry.

  13. 4.2 Electrons • The electron is the smallest subatomic particle with a mass 1/1840 the size of the proton. • An electron has a negative charge (exactly 1-). • English Physicist Sir J.J. Thompson discovered the electron in 1897 using the cathode ray tube. (see p. 86-87)

  14. 4.2 Protons • The proton is a subatomic particle that has a positive charge (exactly 1+). • A proton is a much more massive subatomic particle, when compared to the electron, but it is still very small. • E. Goldstein discovered the proton in 1886 using a cathode ray tube and canal rays (a beam of protons.

  15. 4.2 Neutrons • The neutron is a subatomic particles that have no charge (neutral). • Neutrons have a mass that is about the same as a proton. • An English physicist, Sir James Chadwick, confirmed the existence of the neutron in 1932.

  16. 4.2 Electrons, Protons, Neutrons • The fundamental building blocks of atoms are the electron, the proton, and the neutron. Table 4.1 Properties of Subatomic Particles *1 amu = 1.66x10-24 g

  17. 4.2 Simple Rules about Matter and Electric Charges 1. Atoms have no electric charge – they are electrically neutral. 2. Electric charges are properties of particles. • Electric charges are carried by particles of matter. 3. Electric charges exist in a single unit or multiples of a single unit. • The are no fractions of charges. 4. Electric charges cancel when equal number of positively charged particles are balanced by negatively charged particles.

  18. 4.2 Concept Practice 3. Since all atoms have negatively charged electrons, shouldn’t every sample of matter have a negative charge? Explain. 4. What experimental evidence did Thompson have for the following ideas? a. Electrons have a negative charge. b. Atoms of all elements contain electrons.

  19. 4.3 The Structure of the Nuclear Atom • At first (before the discovery of the neutron) scientists thought that protons and neutrons were evenly distributed throughout the atom.

  20. 4.3 The Structure of the Nuclear Atom • In 1911, Ernest Rutherford and his co-workers tested this theory of atomic structure by firing a beam of alpha particles at a thin gold sheet. • From this experiment Rutherford proposed that almost all of the mass and all of the positive charge are concentrated at a small region at the center of an atom – he called this region the nucleus.

  21. 4.3 The Structure of the Nuclear Atom • nucleus – the central core of an atom, composed of protons and neutrons. • Almost all of mass in an atom is contained in a tiny nucleus which is extremely dense – if it was the size of a pea it would have a mass of 250 tons! • The nucleus has a positive charge and occupies a very small volume of the atom – the rest of the atom is more or less empty space in which the negatively charged electrons are found.

  22. 4.3 Concept Practice 5. How did the results of Rutherford’s gold foil experiment differ from his expectations? 6. What is the charge, positive or negative, on the nucleus of every atom?

  23. 4.4 Atomic Structure Review • Most atoms contain protons (+), neutrons (0), and electrons (-) • Protons and neutrons make up the small, dense nucleus • Electrons surround the nucleus and occupy most of the volume of the atom (mostly empty space) • How are atoms of one element different from atoms of a different element?

  24. 4.4 Atomic Number • Differences among the elements result from differences in the number of protons in an atom. • Example: All atoms of boron (B) have 5 protons, atoms of carbon (C) have 6 protons, and fluorine (F) atoms have 9 protons. atomic number – the number of protons in the nucleus of the atom for a particular element

  25. 4.4 Atomic Number • The atomic number identifies an element • The atomic number for each element is listed on the periodic table (usually above the chemical symbol) • Example: Oxygen always has 8 protons so its atomic number is 8 • Since atoms are electrically neutral – they have the same number of protons (p+) as electrons (e-) – the atomic number indirectly tells the number of electrons

  26. 4.4 Concept Practice 7. Why is an atom electrically neutral? 8. What is the relationship between the number of protons and the atomic number of an atom?

  27. 4.4 Practice 9. Use the periodic table to complete this table.

  28. 4.5 Mass Number • Most of the mass of an atom is concentrated in its nucleus and depends on the number of protons and neutrons mass number – the total number of protons and neutrons in the nucleus mass number = atomic number + number of neutrons or…. number of neutrons = mass number – atomic number

  29. 4.5 Mass Number • You can determine the composition of an atom of any element from its atomic number and mass number. • Example: Oxygen has an atomic number of 8 and a mass number of 16 • Atomic number → 8 protons • Atoms are electrically neutral → 8 electrons • 16 – 8 = 8 neutrons

  30. 4.5 Shorthand Notation • You can represent the composition of any atom with the element’s chemically symbol and two additional numbers written to the left • Example: • When an element name is written with a number after it, the number represent the mass number • Example: gold-197 Au mass number → 197 79 atomic number →

  31. 4.5 Practice 10. Complete the table.

  32. 4.5 Examples • How many neutrons are in the following atoms? a. O b. nitrogen-15 c. Ag 16 8 108 47

  33. 4.5 Concept Practice 11. An atom is identified as platinum-195. a. What is the number 195 called? b. Write the symbol for this atom using superscripts and subscripts.

  34. 4.6 Isotopes of Elements • Most of Dalton’s atomic theory is accepted today, however, it is now known that atoms of the same element may have different nuclear structures. • The nuclei of atoms of a given element must have the same number of protons, but the number of neutrons may vary.

  35. 4.6 Isotopes of Elements isotopes – atoms that have the same number of protons, but different numbers of neutrons • Example: hydrogen-1 → 1 proton, 0 neutrons hydrogen-2 → 1proton, 1 neutron hydrogen-3 → 1 proton, 2 neutrons • Isotopes of an element are chemically alike because they have the same number of protons and electrons • These subatomic particles determine the characteristic chemical behavior of each element

  36. 4.6 Concept Practice 13. How are isotopes of the same element alike? How are they different? 14. Three isotopes of oxygen are oxygen-16, oxygen-17, and oxygen-18. Write the chemical symbol, including the atomic number and atomic mass, for each.

  37. 4.7 Atomic Mass • Even the largest atoms have very small masses (Fluorine – 3.155 x 10-23) • Masses of single atoms can be determined using mass spectrometers, but these masses would be inconvenient to work with. • It is more useful to compare the relative masses of atoms using an isotope of carbon, carbon-12, as a basis.

  38. 4.7 Atomic Mass • Carbon-12 was assigned a mass of exactly 12.00000 amu atomic mass unit (amu) – one-twelfth the mass of a carbon-12 atom • Since carbon-12 has a mass number of 12 (6 protons and 6 neutrons), the mass of a single proton or neutron is approximately 1 amu.

  39. 4.7 Atomic Mass • Since most of the mass of an atom depends on the number of protons and neutrons in the nucleus you may expect the atomic mass to be a whole number – this is not the case.

  40. 4.7 Atomic Mass • In nature, most elements occur as a mixture of two or more isotopes • Each isotope has a fixed mass and a natural percent abundance atomic mass – a weighted average mass of the atoms in a naturally occurring sample of the element • This reflects both the mass and the relative abundance of the isotopes as they occur in nature

  41. 4.7 Atomic Mass • Example: Hydrogen-Atomic mass: 1.0079 amu • Three isotopes: Hydrogen-1, Hydrogen-2, Hydrogen-3 • Hydrogen-1 has a mass of 1.0078 and a natural abundance of 99.98% • Hydrogen-2 and hydrogen-3 occur in trace amounts • The slight difference in the mass of hydrogen-1 and the average atomic mass takes into account the other two isotopes.

  42. 4.8 Calculating the Atomic Mass • Because the atomic mass must reflect both the masses and the relative natural abundances of the isotopes, you must know: • The number of stable isotopes of that element • The mass of each isotope • The natural percent abundance of each isotope

  43. 4.8 Calculating Atomic Mass • Example: Element X has two natural isotopes. The isotope with mass 10.012 amu has a relative abundance of 19.91 %. The isotope with mass 11.009 amu has a relative abundance of 80.09%.Calculate the average atomic mass and determine what element it is.

  44. 4.8 Calculating Atomic Mass • Solution: • Find the mass that each isotope contributes to the weighted average by multiplying the mass by its relative abundance. Then add the products. X 10.012 amu x 0.1991 = 1.993 amu X 11.009 amu x 0.8009 = 8.817 amu Total: 10.810 amu Element X is boron. 10 11

  45. 4.8 Concept Practice 20. The element copper contains the naturally occurring isotopes Cu and Cu. The relative abundances and atomic masses are 69.2% (mass = 62.93 amu) and 30.8% (mass = 64.93 amu), respectively. Calculate the average atomic mass of copper. 63 29 65 29

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