Atoms and Elements Chapter 4

1. Atoms and ElementsChapter 4 Tro, 2nd ed.

2. ELEMENTS AND ATOMS All known substances on Earth and probably the universe are formed by combinations of more than 100 elements. An element is a fundamental or elementary substance that cannot be broken down into simpler substances by chemical means. • Each element has a number. • Beginning with hydrogen as 1, the elements are numbered in order of increasing complexity.

3. ELEMENTS AND ATOMS Most substances can be decomposed into two or more simpler substances. • Water can be decomposed into hydrogen and oxygen. • Table salt can be decomposed into sodium and chlorine. • An element cannot be decomposed into a simpler substance. • An atom is the smallest particle of an element that can exist. • An atom is the smallest unit of an element that can enter into a chemical reaction.

4. An atom is very small

5. The diameter of an atom is 0.1 to 0.5 nm. This is 1 to 5 ten billionths of a meter. Even smaller particles than atoms exist. These are called subatomic particles. If the diameter of this dot is 1 mm, then 10 million hydrogen atoms would form a line across the dot.

6. Modern research has demonstrated that atoms are composed of subatomic particles. Atoms under special circumstances can be decomposed. Dalton’s Atomic Theory • Elements are composed of minute indivisible particles called atoms. • Atoms of the same element are alike in mass and size. Atoms of different elements have different masses and sizes. (Is this still correct?) • Chemical compounds are formed by the union of two or atoms of different elements in whole number ratios.

7. WHAT’S IN AN ATOM? DISCOVERING SUBATOMIC PARTICLES! In 1897 Sir Joseph Thompson demonstrated that cathode rays: - travel in straight lines - are negative in charge - are deflected by electric and magnetic fields - produce sharp shadows - are capable of moving a small paddle wheel

8. Subatomic Particles: Electron This was the discovery of the fundamental unit of charge – the electron.

9. Subatomic Particles: Protons Eugen Goldstein, a German physicist, first observed protons in 1886: Thompson determined the proton’s characteristics. Thompson showed that atoms contained both positive and negative charges. This disproved the Dalton model of the atom which held that atoms were indivisible. But the mass of the atom could not be accounted for by the mass of protons inside it. There had to be something else.

10. Subatomic Particles: Neutron Its actual mass is slightly greater than the mass of a proton. James Chadwick discovered the neutron in 1932.

11. 4.1 0.00055 1.0073 1.0087 Memorize name, symbol, relative charge, relative mass, and LOCATION (in nucleus or in orbitals outside of nucleus) of these three subatomic particles. See table 4.1 on page 90 in your book.

12. The Nuclear Atom Radioactivity was discovered by Becquerel in 1896. Radioactive elements spontaneously emit alpha (a) particles, beta (b) particles and gamma (g) rays from their nuclei. By 1907, Rutherford found that a particles emitted by certain radioactive elements were helium nuclei, consisting of 2 protons and 2 neutrons.

13. The Rutherford Experiment In 1911, Rutherford performed experiments that shot a stream of a particles at a thin piece of gold foil. Most of the a particles passed through the foil with little or no deflection. He found that a few were deflected at large angles and some a particles even bounced back.

14. Rutherford’s alpha particle scattering experiment. 5.5

15. The Rutherford Experiment An electron with a mass of 0.00055 amu could not have deflected an a particle with a mass of ~4 amu. Rutherford knew that like charges repel. Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus.

16. The Rutherford Experiment If a positive a particle approached close enough to the positive mass it was deflected. Most of the a particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space. Because a particles have relatively high masses, the extent of the deflections led Rutherford to conclude that the nucleus was very heavy and dense.

17. Deflection Scattering Deflection and scattering of a particles by positive gold nuclei. 5.5

18. General Arrangement of Subatomic Particles Rutherford’s experiment showed that an atom had a dense, positively charged nucleus. Chadwick’s work in 1932 demonstrated the atom contains neutrons. Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge.

19. General Arrangement of Subatomic Particles Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center. The negative electrons surround the nucleus. The nucleus contains protons and neutrons

20. 5.6

21. ATOMS AND THEIR SUBATOMIC PARTICLES The atomic number of an element is equal to the number ofprotons in the nucleus of that element. The atomic number of an atom determines which element the atom is. Every atom with an atomic number of 1 is a hydrogen atom.Every carbon atom contains 6 protons in its nucleus.

22. atomic number 1 proton in the nucleus 1H Every atom with an atomic number of 1 is a hydrogen atom.

23. atomic number 6 protons in the nucleus 6C Every atom with an atomic number of 6 is a carbon atom.

24. Ions: atoms that have lost or gained electrons Positive ions were explained by assuming that a neutral atom loses electrons. Negative ions were explained by assuming that extra electrons can be added to atoms.

25. When one or more electrons are lost from an atom, a cation is formed. 5.4

26. When one or more electrons are added to a neutral atom, an anion is formed. 5.4

27. Atomic Structures of Ions

28. Isotopes of the Elements Atoms of the same element can have different masses. They always have the same number of protons, but they can have different numbers of neutrons in their nuclei. The difference in the number of neutrons accounts for the difference in mass. These are isotopes of the same element.

29. Isotopic Notation

30. 12 C 6 Isotopic Notation for Carbon-12 6 protons + 6 neutrons 6 protons

31. Isotopic Notation for Carbon-14 6 protons + 8 neutrons 14 C 6 6 protons

32. Hydrogen has three isotopes 1 proton 0 neutrons 1 proton 1 neutron 1 proton 2 neutrons

33. Examples of Isotopes ElementProtonsElectronsNeutronsSymbol Hydrogen 1 1 0 Hydrogen-2 1 1 1 Hydrogen-3 1 1 2 Uranium-235 92 92 143 Uranium-238 92 92 146 Chlorine-35 17 17 18 Chlorine-37 17 17 20

34. Atomic Mass The mass of a single atom is too small to measure on a balance. Using a mass spectrometer, the mass of the hydrogen atom was determined. The mass of one hydrogen atom was determined to be 1.673 x 10-24 g.

35. Positive ions formed from sample. Electrical field at slits accelerates positive ions. Deflection of positive ions occurs at magnetic field. A Modern Mass Spectrometer From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined. A mass spectrogram is recorded. 5.8

36. A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given. 5.9

37. ATOMIC MASS UNITS AND RELATIVE ATOMIC MASS Single atomic masses are too small to weigh on a balance. To overcome this problem a system of relative atomic masses using “atomic mass units” was devised to express the masses of elements using simple numbers. The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon, carbon-12. A mass of exactly 12 atomic mass units (amu) was assigned to the carbon-12 atom. An amu is defined as exactly equal to 1/12th mass of a carbon-12 atom. 1 amu = 1.6606 x 10-24 g Isotopes of the same element have different masses. The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12.

38. ATOMIC MASS UNITS AND RELATIVE ATOMIC MASS Calculating relative atomic mass: Convert percent abundance to fraction by dividing by 100 Multiply fraction abundance times each isotopes mass Sum up the results Rel Atomic Mass = f1*m1 + f2*m2 + … See example on next slide

39. To calculate the atomic mass multiply the atomic mass of each isotope by its fractional abundance and add the results. (62.9998 amu) 0.6909 = 43.48 amu (64.9278 amu) 0.3091 = 20.07 amu 63.55 amu

40. Relationship between Mass Number and Atomic Mass Atomic mass found on the Periodic Table is RELATIVE atomic mass, the weighted average of all the isotopes. Each isotope of an element has a unique WHOLE number called the mass number, the sum of protons and neutrons. Mass number is NOT found on most Periodic Tables. ****************************************************************** ATOMIC MASS IS NOT THE SAME AS MASS NUMBER ATOMIC MASS IS NOT THE SAME AS MASS NUMBER ATOMIC MASS IS NOT THE SAME AS MASS NUMBER ****************************************************************** REMEMBER: ATOMIC MASS IS AN AVERAGE OF THE DIFFERENT ISOTOPES’ MASSES. It will have decimal points!

41. Atomic Mass, Mass Number & Abundance of Isotopes EXAMPLE: Chlorine has two isotopes, chlorine-35 and chlorine-37. Write down the atomic number, mass number and number of neutrons for each. Then look up the atomic mass. Which isotope is in greater abundance? Cl-35: Z = 17, A = 35, 35 - 17 = 18 neutrons Cl-37: Z = 17, A = 37, 37 - 17 = 20 neutrons Atomic mass is 35.453 amu. Closer to 35, therefore chlorine-35 is more abundant. Now do problem 105 in your textbook.

42. Relationship Between Mass Number and Atomic Number atomic number number of neutrons mass number mass number - = atomic number 109 - 47 = The mass number minus the atomic number equals the number of neutrons in the nucleus. 62

43. Oxygen is the most abundant element in the human body (65%). • Oxygen is the most abundant element in the crust of the earth (49.2%). • In the universe, the most abundant element is hydrogen (91%) and the second most abundant element is helium (8.75%). Elements are not distributed equally by nature.

44. Distribution of the common elements in nature. 3.2

45. Sources of Element Names Greek-Color Iodine: from the Greek iodes meaning violet. Latin- Property Fluorine: from the Latin fluere meaning to flow. The fluorine containing ore fluorospar is low melting. German- Color Bismuth: from the German weisse mass which means white mass. Location Germanium: discovered in 1866 by a German chemist. Famous- Scientists Einsteinium: named for Albert Einstein.

46. These symbols have carried over from the earlier names of the elements (usually Latin). A number of symbols appear to have no connection with the element. Most symbols start with the same letter as the element.

48. Element Names and Symbols Element names and symbols are in Periodic Table inside front cover. Symbols: Learn first 36 on Periodic Table, plus Ag, Au, Pt, Hg, Sn, Pb, I, U, Ba, and Rn. Also memorize that some elements exist as diatomic molecules : H2, O2, N2, F2, Cl2, Br2, and I2

49. N THE PERIODIC TABLE OF THE ELEMENTS The periodic table was designed by Dimitri Mendelev in 1869. In the table each element’s symbol is placed inside of a box. Above the symbol of the element is its atomic number. 7

50. He Ne Ar Kr Xe Rn The elements are arranged in order of increasing atomic number. Elements with similar chemical properties are organized in columns called families or groups . These elements are known as the noble gases. They are nonreactive.