1. Unit 3 Chapter 4 Compare the subatomic particles (protons, neutrons, electrons) of an atom with regard to mass, location, and charge, and explain how these particles affect the properties of an atom (including identity, mass, volume, and reactivity). (PS-2.1) Illustrate the fact that the atoms of elements exist as stable or unstable isotopes. (PS-2.2) Use the atomic number and the mass number to calculate the number of protons, neutrons, and/or electrons for a given isotope of an element. (PS-2.4)

2. Early Theories of Matter Early philosophers formed theories based on everyday observations. Thought that there were 4 basic elements: –Earth –Water –Air –Fire 2 forces: love & hate

3. Democritus Greek philosopher (460 – 370 BC) Proposed matter made up of individual particles called atomos –Atoms could not be created, destroyed, or further divided Could not answer question “What holds atoms together?”

4. Aristotle Most influential Greek philosopher (384 – 322 BC) Rejected atomic theory of Democritus because it conflicted with his ideas on nature Did not believe that “nothingness” of empty space could exist

5. John Dalton 19 th century English school teacher (1766 – 1844) Revised Democritus’s ideas based on results of scientific research Proposed his Atomic Theory in 1803 –Explains the Law of Conservation of Mass

6. Defining the Atom The smallest particle of an element that still retains the properties of an element How small is an atom? –World population (2000) 6 x 10 9 people –Atoms in a penny2.9 x 10 22 –Diameter of copper atom = 1.28 x 10 -10 m –6 billion atoms side by side would be less than 1 m long

7. Discovering the Electron Cathode Ray Tube –Tube with metal electrodes on opposite ends –Cathode: electrode connected to negative terminal of battery –Anode: electrode connected to positive terminal of battery Sir William Crookes – found rays traveling from cathode to anode within the tube (Yay! Now we have TV!)

8. Discovering the Electron Research with cathode ray tubes convinced scientists of the following: –Cathode rays were actually a stream of charged particles –The particles carried a negative charge that were found in all forms of matter Electrons: negatively charged particles that are part of all forms of matter

9. J. J. Thomson (1856 – 1940) English physicist Series of cathode ray tube experiments to determine ratio of charge to mass of one cathode ray particle Measured effect of both magnetic & electric fields on cathode ray Determined charge-to-mass ratio of charged particle

10. J. J. Thomson (1856 – 1940) Compared ratio to other known ratios Mass of charged particle much less than hydrogen atom, lightest known atom Conclusion: Dalton’s theory was WRONG! Identified 1 st subatomic particle: electron

11. Robert Milliken (1868 – 1953) American physicist Determined charge of electron (within 1% of currently accepted value) Charge equated to single unit of negative charge (charge e - = -1) Milliken used charge to calculate mass of e - –Mass e - = 9.1 x 10 -28 g = 1/1840 mass H-atom

12. Plum Pudding Model Proposed by J.J. Thompson Reasons: –Matter is neutral so must be something other than electrons –Mass of electron so extremely small so where does the rest of the mass come from? Model: –Uniformly distributed positive charge with electrons inside

14. Ernest Rutherford (1871 – 1937) New Zealand native, won Nobel Prize in chemistry in 1908 1911 – studied how positive alpha particles interacted with matter –Gold Foil Experiment

15. Gold Foil Experiment Shot alpha particles at thin sheet of gold foil Hypothesis: alpha particles would pass straight thru foil with little deflection by small electrons Observations: alpha particles deflected at very large angles, some even straight back toward source Proposed new model to explain observations

16. Nucleus Rutherford proposed: 1.That an atom consisted of mostly empty space through which electrons move 2.Concluded there was a tiny, dense region called nucleus centrally located within atom that contained all of an atom’s positive charge and almost all of its mass 3.electrons move through available space around nucleus & are held within atom by attraction to positive nucleus

17. Protons & Neutrons 1920 – Rutherford revised concept of nucleus –Concluded that nucleus contained positive particles –Proton: subatomic particle carrying charge equal to but opposite that of an electron Charge = +1 1932 – English physicist James Chadwick showed nucleus also contained neutron –Neutron: mass nearly equal to that of proton but has no charge

18. The Atom Neutral particle composed of electrons, protons & neutrons Spherically shaped with tiny, dense, positive center surrounded by 1 or more negative electrons that move quickly through empty space around nucleus Nucleus is 99.97% of atom’s mass Atoms are neutral so # protons = # electrons

19. Atomic Number Henry Moseley (1887 – 1915) –English scientist –Discovered atoms of each element contain a unique positive charge in nucleus # protons in atom identifies it as atom of particular element Atomic number: number of protons in atom (and number of electrons) –Determines element’s place in periodic table

21. Isotopes Atoms of particular element have same number of p & e - but number of neutrons can be different Isotope: atoms with same # of p & e - but different # of neutrons (n) Most elements found as mixture of isotopes –No matter where sample comes from, relative abundance of each isotope is constant Isotopes differ in mass Isotopes of atom have essentially same chemical behavior

22. Mass Number Number added to element’s name to identify isotopes Represents sum of protons and neutrons in nucleus Shorthand notation of isotopes consist of element symbol, mass #, & atomic # # neutrons = mass # - atomic # Examples

23. Mass of Individual Atoms Mass proton = 1.673 x 10 -24 g Mass neutron = 1.675 x 10 -24 g Mass electron = 9.11 x 10 -28 g Chemists developed method of measuring mass of an atom relative to mass of specifically chosen atomic standard – Carbon-12 Assigned carbon-12 atom a mass of exactly 12 atomic mass units

24. Atomic Mass Unit 1 atomic mass unit (amu) = 1/12 th the mass of a carbon-12 atom Mass proton = 1.007 276 amu Mass neutron = 1.008 665 amu Mass electron = 0.000 549 amu

25. Atomic Mass Atomic mass of an element is the weighted average mass of the isotopes of that element Calculated by summing products of each isotope’s percent abundance times its atomic mass

26. Atomic Mass Examples What is the average atomic mass of Boron if it exists as 19.90% 10 B (10.013 g/mol) and 80.10% 11 B (11.009 g/mol)?

27. Atomic Mass Examples What is average atomic mass of Lithium if 7.42% exists as 6 Li (6.015 g/mol) and 92.58% exists as 7 Li (7.016 g/mol)?

28. Atomic Mass Examples Calculate the atomic mass of magnesium. The three magnesium isotopes have atomic masses and relative abundances of 23.985 amu (78.99%), 24.986 amu (10.00%), and 25.982 amu (11.01%).

30. Bohr Model of the Atom Niels Bohr –Danish physicist –Worked for Rutherford in 1913 –Proposed quantum model of atom that seemed to explain the discontinuous spectra of elements –Model correctly predicted frequencies of lines in hydrogen’s atomic emission spectrum

31. Bohr Model of the Atom Energy states of hydrogen –Proposed certain allowable energy states for H- atom –Ground state: lowest allowable energy state of atom –Excited state: when atom gains energy –Related H-atom energy states to motion of electrons within atom electrons moved around nucleus in certain allowed circular orbits

32. Bohr Model of the Atom Orbitals –Quantum number, n, assigned to each orbit –Calculated orbit radius

33. Bohr Model of the Atom Hydrogen’s line spectrum –When electron of excited H-atom drops from high energy level to lower, photon emitted –Orbits are like rung of ladder – 7 orbits in H-atom Model did not work on any other atom!

34. Quantum Mechanical Model 1924 – Louis de Broglie (1892 – 1987) –French graduate student in physics –Compared Bohr’s orbitals to waves – only whole numbers of wavelengths allowed in circular orbit of fixed radius –Question: Can particles of matter, including electrons, behave like waves? –Answer: If electron has wavelike motion & is restricted to circular orbits of fixed radius, the electron is allowed only certain possible wavelengths, frequencies, and energies

35. Heisenberg Uncertainty Principle Werner Heisenberg (1901 – 1976) –German theoretical physicist Impossible to make any measurement on an object without disturbing the object Fundamentally impossible to know precisely both the momentum and position of a particle at the same time

36. Schrödinger Wave Equation Erwin Shrödinger (1887 – 1961) –Austrian physicist –1926 – further work of de Broglie Derived equation treating hydrogen’s electron as a wave Equation applied to atoms of other elements Limits electron’s energy to certain values (like Bohr) No attempt to describe electron’s path around nucleus

37. Schrödinger Wave Equation Solution called wave function Atomic orbital: 3-D region around nucleus (electron cloud) –Density of cloud at given point proportional to probability of finding electron at that point

38. Hydrogen’s Atomic Orbitals Principal quantum number (n) –Indicates relative sizes and energies of atomic orbitals –As n increases: orbital becomes larger Electron spends more time farther from nucleus Atom’s energy level increases –Specifies atom’s major energy levels, principal energy levels –For hydrogen atom, 1 ≤ n ≤ 7

39. Hydrogen’s Atomic Orbitals Energy sublevels –Contained in principal energy levels –Number of energy sublevels in a principal energy level increases as n increases –Labeled s, p, d, f according to shape of orbital s orbital – spherical p orbital – dumbbell shaped d and f orbitals not always same shape Each orbital can only contain 2 electrons

40. Atomic Emission Spectra Set of frequencies of electromagnetic (light) waves emitted by atoms of an element Similar to fingerprint –Each element has unique emission spectra –Spectra can be used to identify elements –Only certain colors appear in spectrum therefore, only photons having certain specific energies are emitted by excited atoms