Chapter 4

1. Chapter 4 Elements, Atoms, and Ions

2. Ch#4 Contents 4.1 The Elements 4.2 Symbols for the Elements 4.3 Dalton’s Atomic Theory 4.4 Formulas of Compounds 4.5 The Structure of the Atom 4.6 Modern Concept of Atomic Structure 4.7 Isotopes 4.8 Introduction to the Periodic Table 4.9 Natural States of the Elements 4.10 Ions 4.11 Compounds that Contain Ions 4.9 Natural State of the Elements 4.10 Ions 4.11 Compounds that Contain Ions

3. Elements A substance that cannot be broken down by chemical means. Elements are defined by the number of protons they possess. • 115 known: 88 found in nature, others are man made. • Most of the know elements are found on the Periodic Chart handing on the wall. • One or two letters are used to represent an element and they are called the symbol • Some elements are solids (black), some are liquids (blue) and some are gases (red)

4. Element Distribution in Earth’s Surface

5. Symbols • Each element has a unique one- or two-letter symbol. • First letter is always capitalized and the second is not. • The symbol usually consists of the first one or two letters of the element’s name. • Examples: Oxygen O Krypton Kr • Sometimes the symbol is taken from the element’s original Latin, Greek, English, or German name. • Examples: Gold Au aurum Lead Pbplumbum Wolfram W tungsten

6. Symbols Located on the front page of your textbook

7. Dalton’s Atomic Theory In 1808 Dalton published A New System of Chemical Philosophy, where he presented his theory of atoms: • Each element is made up of tiny particles called atoms. • The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways • Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms. • Chemical changes involve reorganization of the atoms, to different ratios.

8. Dalton’s Atomic Theory Which of the following statements regarding Dalton’s atomic theory are still believed to be true? • Elements are made of tiny particles called atoms. • All atoms of a given element are identical. • A given compound always has the same relative numbers and types of atoms. • Atoms are indestructible.

9. Dalton’s Atomic Theory Which of the following statements regarding Dalton’s atomic theory are still believed to be true? • Elements are made of tiny particles called atoms. • All atoms of a given element are identical. • A given compound always has the same relative numbers and types of atoms. • Atoms are indestructible.

10. Mendeleev’s Periodic Table • Dmitrii I. Mendeleev arranged elements in the periodic table by their chemical and physical properties. • He left open spaces in his periodic table to account for elements not yet discovered.

11. The Modern Periodic Table • The modern periodic table is also based on a classification of elements in terms of their physical and chemical properties. • The horizontal rows are called periods. • Columns contain elements of the same family or group. • Transition metals are the elements in group 3 through 12 in the periodic table.

12. Formulas • Compound – distinct substance that is composed of the atoms of two or more elements and always contains exactly the same whole number ratio of those elements. • Chemical Formulas – expresses the types of atoms and the number of each type in each formula unit, or molecule of a given compound. • Formula Examples • H2O (molecule) • C6H12O6 (molecule) • NaCl (Formula Unit)

13. Metals

14. Atomic Structure J. J. Thomson (1898—1903) • Postulated the existence of electrons using cathode-ray tubes. • The atom must also contain positive particles that balance exactly the negative charge carried by particles that we now call electrons.

15. Cathode-ray Tube

16. Atomic Structure William Thomson (Plum Pudding Model) Reasoned that the atom might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to counterbalance that positive charge.

17. Atomic Structure Ernest Rutherford (1911) • Explained the nuclear atom. • Atom has a dense center of positive charge called the nucleus. • Electrons travel around the nucleus at a relatively large distance. • A proton has the same magnitude of charge as the electron, but its charge is positive.

18. Atomic Structure Rutherford and Chadwick (1932) • Most nuclei also contain a neutral particle called the neutron. • A neutron is slightly more massive than a proton but has no charge.

19. Atomic Structure • Electrons – found outside the nucleus; negatively charged • Protons– found in the nucleus; positive charge equal in magnitude to the electron’s negative charge • Neutrons– found in the nucleus; no charge; virtually same mass as a proton

20. Atomic Structure • The nucleus is: • Small compared with the overall size of the atom. • Extremely dense; accounts for almost all of the atom’s mass.

21. Atomic Structure

22. Atomic Structure Why do different atoms have different chemical properties? • The chemistry of an atom arises from its electrons. • Electrons are the parts of atoms that “intermingle” when atoms combine to form molecules. • It is the number of electrons that really determines chemical behavior.

23. Isotopes • Atoms with the same number of protons but different numbers of neutrons. • Show almost identical chemical properties; chemistry of atom is due to its electrons. • In nature most elements contain mixtures of isotopes.

24. Symbols of Isotopes • Each element consists of atoms with the same number of protons in the nucleus. This number is called atomic number (Z). • Protons and neutrons in atomic nuclei are called nucleons. • The mass number (A) is the total number nucleons in one atom of an element.

25. A X Z Writing the Symbol of an Isotope • A is the mass number • Z is the atomic number • X is the atomic symbol

26. Isotopes • C = the symbol for carbon • 6 = the atomic number (6 protons) • 14 = the mass number (6 protons and 8 neutrons) • C = the symbol for carbon • 6 = the atomic number (6 protons) • 12 = the mass number (6 protons and 6 neutrons)

27. Two Isotopes of Sodium

28. BOHR MODELS OF HYDROGEN (The three isotopes of hydrogen) E E E N P P N N P N P N N Hydrogen-3 Hydrogen-2 Hydrogen-1

29. BOHR MODELS OF HYDROGEN E E E P P P N N N N N P N Hydrogen-3 Hydrogen-2 Hydrogen-1 2 3 1 H H H

30. BOHR MODELS OF HYDROGEN E E E P P P N N N N N P N Hydrogen-3 Hydrogen-2 Hydrogen-1 1 2 3 H H H 1 1 1

31. BOHR MODELS OF HYDROGEN E E E P P P N N N N N P N Hydrogen-3 Hydrogen-2 Hydrogen-1 1 2 3 H H H 1 1 1 Atomic Number (number of protons); found on the periodic chart

32. BOHR MODELS OF HYDROGEN E E E P P P N N N N N P N Hydrogen-3 Hydrogen-2 Hydrogen-1 1 2 3 H H H 1 1 1 Atomic Number (number of protons) Mass number (sum of protons and neutrons); Not found on periodic chart

33. BOHR MODELS OF HYDROGEN E E E P P P N N N N N P N Hydrogen-3 Hydrogen-2 Hydrogen-1 1 2 3 H H H 1 1 1 Atomic Number (number of protons) Mass number (sum of protons and neutrons); Not found on periodic chart Oxidation number (Protons – electrons)

34. Relative Atomic Mass Some atoms naturally occur as a mixture of isotopes. e.g. 11H, 12H and 13H The atomic weights given in the periodic table take this into account. They are the average atomic weight taking into account the amounts of each isotope present.

35. Relative Atomic Mass of Hydrogen When computing the average atomic mass of any element the radioactive (unstable) isotopes are excluded since there relative abundances are slowly decreasing. Hydrogen has three isotopes. The first two protium and deuterium are stable isotopes and the third tritium is unstable, thus excluded in the calculation. The relative abundance of isotopes of a particular element is constant here on our planet Earth. Mass spectroscopy gives information of relative abundances and relative masses of isotopes.

36. Relative Atomic Mass of Hydrogen Mass spectroscopy gives the following information for the two stable isotopes of hydrogen. isotope namerelative abundanceMass (amu) Hydrogen-1 (protium) 99.985 % 1.007825 Hydrogen-2 (deuterium) 2.01355 0.015 % To calculate a weighted average, convert the percent to a decimal by moving the decimal two places to the left. Then multiply this number by the mass and add the two masses.

37. Relative Atomic Mass of Hydrogen 0.99985 X 1.007825 = 1.007673 0.0015 X 2.01355 = 0.003021 1.010694 amu 1.0107 amu

38. Average Atomic Masses • A mass spectrometer is an instrument that measures precise masses and relative amounts of ions of atoms and molecule. • The natural abundance of an isotope is its relative proportion among all the isotopes found a natural sample. • The average atomic mass of an element is calculated by multiplying the natural abundance of each isotope by its exact mass in atomic mass units and then summing these products.

39. Example Neon is found in three isotopes in nature. 19.9924 x 0.904838 = 18.0898323 amu 20.99395 x 0.002696 = 0.056599689 amu 21.9914 x 0.092465 = 2.033434801 amu 20.1797 amu

40. Relative Atomic Mass If we assign a mass of 12 atomic mass units (u or sometimes amu) to a carbon-12 atom then we can compute the relative atomic weight for any other atom. Atomic weights or atomic masses are given in your periodictable.

41. Natural State of Elements • Most elements are very reactive. • Elements are not generally found in uncombined form. • Exceptions are: • Noble metals – gold, platinum and silver • Noble gases – Group 8

42. Diatomic Molecules

43. Allotropes • Different forms of a given element. • Example: • Solid carbon occurs in three forms. • Diamond • Graphite • Buckminsterfullerene

44. Carbon Allotropes

45. Ions • Atoms can form ions by gaining or losing electrons. • Metals tend to lose one or more electrons to form positive ions called cations and are named by using the name of the parent atom. • Nonmetals tend to gain electrons to form negative ions called anions and are named by using the root of the atom name followed by the suffix –ide. • Atoms become more stable by gaining or losing electrons. • Cations are generally named by using the name of the parent atom.

46. Common Oxidation States

47. Common Oxidation States 1+ 2+ 1-

48. Common Oxidation States By Group Number

49. Sample Problem An ion with a 3+ charge contains 23 electrons. Which ion is it?

50. Sample Problem An ion with a 3+ charge contains 23 electrons. Which ion is it? Fe3+ A certain ion X+ contains 54 electrons and 78 neutrons. What is the mass number of this ion?