1. CHAPTER 4 AtomicStructure

2. Democritus (4 th Century B.C.) ► First suggested the existence of tiny particles called atoms (atomos) ► Atoms were indivisible and indestructible ► Not useful for explaining chemical behavior ► Lacked experimental support (no scientific testing)

3. John Dalton (1766-1844) ► Performed tests and experiments ► Studied the ratios in which elements combine in reactions ► Formulated hypothesis & theory

4. Dalton’s Atomic Theory 1. All elements are composed of tiny particles called atoms. Atoms are indivisible and indestructible particles. 2. Atoms of the same element are identical. 3. Atoms of different elements are different. 4. Compounds are formed by the joining of atoms of two or more elements.

5. ATOM ► Smallest particle of an element that retains the properties of that element ► Can be moved & arranged in patterns (scanning tunneling microscope allows us to see this) ► 100,000,000 Cu atoms aligned approximately equals 1cm

6. Revisions to Dalton’s Atomic Theory ► atoms can be broken down into smaller particles: protons, electrons, and neutrons ► Due to the discovery of isotopes, atoms may have different masses. ► Nuclear changes can cause changes in atoms.

7. J.J. Thomson (1897) ► Electrons- negatively charged particles ► Passed electric current through gases at low pressure ► Sealed gases in glass tubes with electrodes ► Electrodes connected to high-voltage electricity

8. ► Anode- positively charged electrode ► Cathode- negatively charged electrode ► Cathode ray- glowing beam travels from cathode to anode ► Cathode rays are attracted to anode and repelled by cathode ► Opposites attract  Thomson proposed that a cathode ray is a stream of tiny negatively charged particles moving at high speed

9. Thomson’s Conclusions ► Electrons are negatively charged particles ► Production of cathode ray does not depend on gas or metal used ► Electrons are part of atoms of all elements ► Electron mass = 1/2000 mass H atom

10. Matter and Electric Charges ► Atoms have no electrical charge ► Electric charges are carried by particles of matter ► Electric charges are always whole numbers ► Negatively charged particles combine with an equal number of positively charged particles to yield a neutral particle

11. Rutherford (1911) ► Tested theory of atomic structure ► Directed narrow beam of alpha particles at very thin sheet of Au foil ► Most alpha particles passed through the Au atoms without deflection ► Small fraction bounced off Au foil at large angle ► Some bounced back toward the source

12. Rutherford’s Conclusions ► Proposed that atom is mostly empty space ► All positive charge & most of the mass are concentrated in small region called the nucleus ► Nucleus- central core of the atom & is composed of protons and neutrons

13. Bohr Model ► Revised Rutherford’s model ► Electrons move in definite orbits around the nucleus ► Orbits = energy levels. ► Energy levels are certain distances from the nucleus ► Farther away from the nucleus the higher the energy

14. Wave Model ► Based on Rutherford’s and Bohr’s model ► Uses the principles of wave mechanics ► Involve complex mathematical equations ► Electron’s do not move about in definite paths ► It is impossible to determine the exact location of an electron. ► Can only predict where an electron can be found ► The probable location is based on the amount of energy the electron has. ► An atom has a small positively charged nucleus surrounded by a large region in which there are enough electrons to make the atom neutral.

15. Atomic Number ► Elements are different because they contain different numbers of protons. ► Example: H atom – 1 proton O atom- 8 protons ► The number of protons in the nucleus of an atom of an element ► Identifies the element ► Atoms are neutral  Number of protons equals the number of electrons ► The atomic number is equal to the number of electrons

16. Mass Number ► Mass number equals the total number of protons plus total number of neutrons ► Example: He atom 2 protons, 2 neutrons Mass number He atom = 4 ► C atom 6 protons, 6 neutrons Mass number C atom = 12

17. ► Atomic number & mass number of an atom of any element can be used to determine the atom’s composition ► Number of neutrons=mass # - atomic # ► Example: O atom atomic # = 8 O atom mass # =16 16 – 8 = 8 neutrons Determining # of Neutrons

18. Shorthand Notation 197 79 Au 79 Au Mass # = 197 Atomic # = 79 How many neutrons does an Au atom have? # neutrons = 197 – 79 = 118 = 118

19. Examples ► Be  Atomic # 4, Mass # 9 9 - 4 = 5 neutrons ► Ne  Atomic # 10, Mass # 20 20 – 10 = 10 neutrons ► Na  Atomic # 11, Mass # 23 23 – 11 = 12 neutrons

20. Isotopes ► Atoms that have the same number of protons but different numbers of neutrons ► Isotopes of elements have different mass numbers ► Chemically alike because they have the same number of protons and electrons Protons and electrons are responsible for chemical behavior

21. Isotopes of Hydrogen ► Hydrogen 1 proton, 0 neutrons Mass # = 1 ► Deuterium 1 proton, 1 neutron Mass # = 2 ► Tritium 1 proton, 2 neutrons Mass # = 3

22. ► Mass spectrometer – used to determine the mass of protons, neutrons, electrons Mass of proton =1.67 * 10 -24 g Mass of neutron = 1.67 * 10 -24 g Mass of electron = 9.11 * 10 -28 g ► More useful to compare relative masses of atoms using a reference isotope as a standard ► Chosen isotope is C-12 has a mass of exactly 12 atomic mass units (amu)

23. Atomic Mass Unit (amu) ► 1/12 th the mass of a C-12 atom ► He-4 atom mass = 4.0026 amu About 1/3 the mass of a C-12 atom ► C-12 atom has 6 protons & 6 neutrons in nucleus  mass of a single proton or neutron is equal to 1 amu

24. ► The mass of any single atom depends mainly on the # of protons and neutrons in the nucleus Prediction  atomic mass of an element is a whole number ► Need to consider the relative abundance of naturally occurring isotopes of the element ► Most elements occur in nature as a mixture of 2 or more isotopes ► Each isotope has a fixed mass & percent abundance

25. Example 2 stable isotopes of chlorine Cl 35 & Cl 37 (34.969 amu + 36.966 amu)/2 = 35.968 amu This is higher than the actual value because Cl 35 occurs 75% naturally Cl 37 occurs 25% naturally Cl 37 occurs 25% naturally Actual atomic mass = 35.453 amu

26. Atomic Mass ► Weighted average mass of the atoms in a naturally occurring sample of the element ► Reflects both the mass & relative abundance of the isotopes ► Calculate atomic mass based on: # of stable isotopes of the element mass of each isotope natural percent abundance of each isotope

27. Example Element X has 2 natural isotopes Isotope 10.012 amu ( 10 X) 19.91 % Isotope 11.009 amu ( 11 X) 80.09% Mass each isotope * relative abundance 10.012 amu * 0.1991 = 1.993amu 11.009 amu * 0.8009 = 8.817 amu 10.810 amu 10.810 amu

28. Forces Within the Atom ► There are four forces that govern all the interactions between matter and energy: ► Electromagnetic Force ► Strong Force ► Weak Force ► Gravity

29. Electromagnetic Force ► Attracts or repels the particles on which it acts. ► Opposite charges attract (protons and electrons) ► Like charges repel (protons) ► Electrons are kept in orbit around the nucleus by the electromagnetic force.

30. Strong Force ► Opposes the electromagnetic force of repulsion between protons. ► Helps to form the nucleus ► Greatest of the 4 forces ► Keeps the protons from repelling each other

31. Weak Force ► Responsible for the process of radioactive decay ► During the process of radioactive decay a neutron in the nucleus changes into a proton and an electron.

32. Gravity ► Weakest of the 4 forces ► Is the force of attraction between all objects known in nature. ► Effects of gravity are easy to observe with large objects ► It is not yet understood how gravity effects the nucleus of an atom.