Chapter 4 Atomic Structure

1. Chapter 4Atomic Structure

2. The Atom You cannot see the tiny fundamental particles that make up matter. Yet, all matter is composed of such particles, called atoms Atom – the smallest particles of an element that retains its identity in a chemical reaction Several early philosophers and scientists could not observe individual atoms, but still were able to propose ideas on the structure of atoms.

3. Democritus’s Atomic Philosophy • Greek philospher Democritus (460B.C – 370 B.C.) was among the first to suggest the existence of atoms. • Democritus believed that matter consisted of tiny, indivisible and indestructible. • Democritus’s ideas did not explain chemical behavior. • Lacked experimental support, because his approach was not based on scientific method.

4. Dalton’s Atomic Theory The modern process of discovery regarding atoms began with John Dalton, an English chemist and school teacher. Dalton used experimental methods and transformed Democritus’s ideas on atoms into scientific theory. Dalton studied the ratios in which elements combine in chemical reactions. Based on the results of his experiments, Dalton formulated hypotheses and theories to explain his observations.

5. Dalton’s Atomic Theory • According to Dalton’s atomic theory, and element is composed of only one kind of atom, and a compound is composed of particles that are chemical combinations of different kinds of atoms. • All elements are composed of tiny indivisible particles called atoms • Atoms of the same element are identical. The atoms of any one element are different from those of any other element.

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

7. The Atom The radii of most atoms fall within the range of 5 x 10-11 m to 2 x 10-10m. (very small) Individual atoms are visible with instruments such as scanning tunneling microscopes.

8. End of section 4.1

9. Subatomic Particles • Most of Dalton’s atomic theory is accepted today. Except, we now know atoms to be divisible. • Atoms can be broken down into smaller particles, called subatomic particles. • There are 3 kinds of subatomic particles. • electrons • Protons • neutrons

10. Electrons In 1897, English physicist J.J. Thomson discovered the electron. Electrons – negatively charged subatomic particles. Dalton performed experiments that involved passing electric current through gases at low pressure.

11. Cathode-Ray Tube Sealed gases in glass tubes fitted at both end with metal disks called electrodes. Electrodes were connected to a source of electricity. One electrode (anode) became positively charged. The other electrode (cathode) became negatively charged. Result was a glowing beam (cathode ray) that traveled from the cathode to the anode.

12. Cathode-Ray Tube Positively charged metal plate attracts the cathode ray, while a negatively charged plate repels it. Thomson knew that opposite charges attract and like charges repel. He hypothesized that a cathode ray is a stream of tiny negatively charged particles (electrons) moving at high speed.

13. Cathode-Ray Tube To test his hypothesis, Thomson set up an experiment to measure the ratio of the charge of an electron to its mass. He found this ratio to be constant. In addition, the charge-to-mass ratio of electrons did not depend on the kind of gas in the tube or the type of metal used for the electrodes. He concluded that electrons must be parts of the atoms of all elements.

14. The Electron An electron carries exactly one unit of negative charge The electrons mass is 1/1840 the mass of a hydrogen atom. How do negatively charged plates affect the path of cathode rays? The negatively charged plate repels the cathode ray.

15. Protons and Neutrons • After a hydrogen atom loses an electron, what is left? • Atoms have no net electric charge, they are electrically neutral • Electric charges are carried by particles of matter • Electric charges always exist in whole-number multiples of a single basic unit. • When a given number of negatively charged particles combines with an equal number of positively charged particles, and electrically neutral particle is formed.

16. Protons and Neutrons • After a hydrogen atom loses an electron, what is left? • A particle with one unit of positive charge should remain when a typical hydrogen atom loses an electron. • In 1886, Eugene Goldstein observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. • He concluded they were positive particles. • Protons – positively charged subatomic particles.

17. Protons and Neutrons English physicist James Chadwick confirmed the existence of another subatomic particle. Neutron – subatomic particles with no charge but with a mass nearly equal to that of a proton.

18. Atomic Nucleus J.J. Thomson though that electrons were evenly distributed throughout an atom filled uniformly with positively charged material. Electrons stuck into a lump of positive charge. This model of the atom was short-lived due to work of Ernest Rutherford, a former student of Thomson.

19. Rutherford’s Gold-Foil Experiment In 1911, Rutherford used alpha particles (Helium atoms that have lost their two electrons and have a double positive charge because of the two remaining protons) to test the current theory of atomic structure. The experiment used a narrow beam of alpha particles directed at a very thin sheet of fold foil. The alpha particles should have passed easily through the gold with only slight deflection.

20. Rutherford’s Gold-foil Experiment However, the great majority of alpha particles passed straight through the gold atoms, without deflection. Also, a small fraction of the alpha particles bounced off the gold foil at very large angles.

21. Rutherford’s Gold-foil Experiment Based on his experimental results, Rutherford suggested a new theory of the atom. He proposed that the atom is mostly empty space, thus explaining the lack of deflections of most of the alpha particles. He concluded that all the positive charge and almost all the mass are concentrated in a small region that has enough positive charge to account for the great deflection . Nucleus – the tiny central core of an atom and is composed of protons and neutrons.

22. Rutherford’s Atomic Model Rutherford atomic model is know as the nuclear atom In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom. The nucleus is tiny compared with the atom as a whole. Although an improvement over Thomson’s model of the atom, Rutherford’s model turned out to be incomplete and had to be modified (chapter 5)

23. Questions What are 3 types of subatomic particles? Proton, neutron, & electrons. How does the Rutherford model describe the structure of atoms? A positively charged nucleus surrounded by electrons, which occupy most of the volume.

24. Questions Describe Thomson’s and Millikan’s contributions to atomic theory. Thomson – Cathode ray experiments which concluded that electrons must be parts of the atoms of all elements. Millikan determined the charge and mass of the electron. What experimental evidence led Rutherford to conclude that an atom is mostly empty space? The great majority of the alpha particles passed straight through the gold foil

25. Questions Compare Rutherford’s expected outcome of the gold-foil experiment with the actual outcome. Expected all alpha particles to pass straight through with little deflection. Found that most passed straight through, but some particles were deflected at large angles and some bounced back.

26. End of Section 4.2

27. Distinguishing Among Atoms How are atoms of hydrogen different from atoms of oxygen? Elements are different because they contain different number of protons. Atomic number – of an element is the number of protons in the nucleus of an atom of that element. Example – all hydrogen atoms have 1 proton and the atomic number of hydrogen is 1. The atomic number identifies an element.

28. Distinguishing Among Atoms 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 an atom Example: Helium atom contains 2 protons and two neutrons, so its mass number is 4 If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition.

29. Distinguishing Among Atoms Example: Oxygen Atomic number is 8 = number of p+ = e- (So oxygen has 8 electron s and 8 protons.) Mass number is 16 = number of p+ plus the number of n0. (So oxygen has 8 neutrons) Number of neutron = mass number – atomic number 197 Au 79 Mass number Atomic number

30. Isotopes There are some elements that have different kinds of atoms of the same element Example – there are three different kinds of Neon atoms Isotopes – are atoms that have the same number of protons, but different numbers of neutrons. Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. Isotopes are chemically alike because they have identical numbers of protons and electrons, which are the subatomic particles responsible for chemical behavior.

31. Hydrogen Isotopes 0 neutrons Mass # - 1 1 neutron Mass # - 2 2 neutrons Mass # - 3

32. Chemical Symbols of Isotopes Write the chemical symbols for three isotopes of oxygen. Oxygen 16, oxygen 17, and oxygen 18. Mass Number (# protons + # neutrons) 16 17 18 O O O 8 8 8 Atomic number (# proton = # electrons)

33. Atomic Mass Actual masses of individual atoms are small and impractical to work with. It is more useful to compare the relative masses of atoms using a reference isotope as a standard The carbon-12 atom was assigned a mass of exactly 12 atomic mass units. Atomic mass unit (amu) – one twelfth of the mass of carbon-12 atom.

34. Atomic Mass In nature, most elements occur as a mixture of two or more isotopes. Each isotope of an element has a fixed mass and a natural percent abundance. Example – almost all naturally occurring hydrogen (99.9985%) is hydrogen-1. The other two isotopes are present in trace amounts. The atomic mass of hydrogen is 1.0079 amu, and Is very close to the mass of hydrogen-1 (1.0078 amu)

35. Atomic Mass The slight difference takes into account the larger masses, but smaller amounts of the other two isotopes of hydrogen. Atomic mass – of an element is a weighted average mass of the atoms in a naturally occurring sample of the element. The atomic mass of copper is 63.546 amu. Which of copper’s two isotopes is more abundant: copper -63 or copper-65? Atomic mass of 63.546 is closer to 63 than 65, thus copper-63 must be more abundant.

36. Atomic Mass Atomic mass = multiply the mass of each isotope by its natural abundance, expresses as a decimal, and then add the products. Element X has two natural isotopes. The isotope with a mass of 10.012 amu has a relative abundance of 19.91%. The isotope with a mass of 11.009 amu has a relative abundance of 80.09%. Calculate the atomic mass of this element. (10.012 amu x 0.1991) + (11.009 amu x 0.8009) (1.993 amu) + (8.817 amu) Atomic mass = 10.810

37. Question Copper – 63 has a mass of 62.93 amu and 69.2% abundance. Copper-65 has a mass of 64.93 amu and 30.8% abundance. What is copper’s average atomic mass? (62.93 amu x 0.692) + (64.93 amu x 0.308) (43.548 amu) + (19.998 amu) Atomic mass = 63.55

38. Periodic Table Periodic Table – an arrangement of elements in which the elements are separated into groups based on a set of repeating properties.

39. Periodic Table Each element is identified by its symbol place in a square. The atomic number of the element is shown centered above the symbol. Elements are listed in order of increasing atomic number, from left to right and from top to bottom. Period - each horizontal row of the periodic table. Within a given period, the properties of the elements vary as you move across it from element to element. Group – each vertical column of the periodic table. Elements within a group have similar chemical and physical properties. Each group is identified by a number and the letter A or B.

40. Periodic Table What distinguishes the atoms from one element from the atoms of another? The number of protons What equation tells you how to calculate the number of neutrons in an atom? Mass number – atomic number = # of neutrons. How do the isotopes of a given element differ from one another? Different mass number and different numbers of neutrons.

41. Periodic Table What makes the periodic table such a useful tool? It allows you to compare the properties of the elements What does the number represent in the isotope platinum-194? Write the symbol from this atom using superscripts and subscripts. It represents the mass number 194 Pt 78

42. Periodic Table Name the elements that have properties similar to those of the element calcium (Ca). Beryllium (Be), magnesium (Mg), strontium (Sr), Barium (Ba), radium (Ra) 194 Consider Pt how would changing the value of the 78 subscript change the chemical properties of the atom? The subscript is the number of protons in atoms of the isotope. Changing the number of protons would change the chemical identity of the isotope to that of another element.

43. End of Chapter 4