Which of these things is not like the other?

1. Which of these things is not like the other? Made By: Ms. Towle

2. It’s All Relative is to as is to

3. Evolution of Atomic Theory – 400BC Democritus 400 BC defined a particle of matter as an Atom (Greek for indivisible).

4. Late 18th Century • Element • Cannot be broken down by ordinary chemical means. • Combine to form compounds that have different chemical and physical properties than those of the individual elements that formed them.

5. Late 18th Century • Chemical Reaction • The transformation of a substance or substances into one or more new substances. 2H2 + O2 2H2O

6. Late 18th Century • Law of Conservation of Matter • Mass is neither destroyed nor created during ordinary chemical reactions or physical changes . 4.0 g H2 + 64.0 g O2 68.0 g H2O

7. Late 18th Century • Law of Definite Proportions • A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. A water molecule is always H2O

8. Late 18th Century • Law of Multiple Proportions • If two or more different compounds are composed of the same elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers. Water = H2O Hydrogen Peroxide = H2O2. Ratio of Oxygen is 1:2 (H2 is the fixed mass)

9. Dalton’s Atomic Theory • All matter is composed of extremely small particles called atoms. Atoms cannot be subdivided, created, or destroyed. • Atoms of the same element are identical in size, mass, and other properties: atoms of different elements differ in size, mass, and other properties. • Atoms of different elements combine in simple whole-number ratios to form chemical compounds. • In chemical reactions, atoms are combined, separated, or rearranged.

10. Errors in Dalton’s Theory 1. We now know that atoms of an element can differ in mass (isotopes). 2. We now know that atoms are made of smaller particles. (Subatomic Particles).

11. J.J. Thomson and the Electron • Cathode Rays and Electrons • Late 1800’s, Joseph John (J.J.) Thomson conducted a series of experiments to investigate the hypothesis that the particles were negatively charged.

12. J.J. Thomson and the Electron • He measured ratio of the particles’ charge to mass and it was the same regardless of variable gases (inside tube) or metals (cathode). Plum Pudding Model

13. Rutherford and the Nucleus • 1911, Ernest Rutherford conducted an experiment in which he bombarded a thin gold foil with alpha particles (positively charged, same as He nucleus).

14. Rutherford and the Nucleus • Hypothesis: Since gold atoms have uniform mass and charge the particles should pass right through. • Results and Conclusion: Most did, however some (1:8,000) had been driven back to the source. • 2 years later, he concluded a positively charged “bundle of matter” occupying a small amount of space must be present in the atom. He named it the nucleus

15. Chadwick and the Neutron • 1932 James Chadwick • Rutherford had proposed a neutral particle to account for the difference in Mass and Atomic numbers. • Bombarded Be with alpha particles, recorded results of particles emitted. • Found a neutral component with the same mass as a proton. • He called it the Neutron.

16. From Rutherford to Bohr • Investigations with light and inquiry is to how electrons were arranged around the nucleus led to the Bohr model of the atom. • Visible light is a kind of electromagnetic radiation from the sun. • Electromagnetic radiation is energy that exhibits wavelike behavior. • Visible light ranges from 400nm (violet) to 700nm(red) in wavelength.

17. From Rutherford to Bohr • Visible light travels at the same velocity regardless of its wavelength. • This velocity is the speed of light: 3.0 x 108 m/s. • The relationship between the wavelength and the frequency of light is c = λf. • Therefore, a shorter wavelength corresponds to a higher frequency.

18. From Rutherford to Bohr • In the early 1900’s, Max Planck theorized another description of light. • He explained from his study on the emission of light from objects that the light was not emitted in continuous waves but in small, specific amounts of energy called quanta. • A quantum is the minimum amount of energy that can be gained or lost by an atom calculated with E = hf.

19. From Rutherford to Bohr • A few years later, Einstein proposed that both theories were correct and light energy has both wave and particle nature. • Light energy can also be a stream of particles. • The particles are called photons. • A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.

20. Rutherford to Bohr • Scientists were also investigating the potential energy of atoms when a current was passed through a gas at low pressure. • The lowest potential energy state of an atom is known as its ground state. • Any higher potential energy state than the energy in an atom’s ground state is known as an atom’s excited state.

21. From Rutherford to Bohr • When a narrow beam of emitted light, produced when the current was applied to the gas, was passed through a prism a line emission spectrum was produced. • The line emission spectrum of a hydrogen atom is: • A photon corresponds to each frequency in the line emission spectrum. The energy of the photon is equal to the difference between the two energy states of an atom.

22. Bohr’s Model of the Atom • In 1913 Bohr theorized a way to explain the line emission spectrum of hydrogen. • He theorized that it was hydrogen’s electron that was related to the photon emission. • He proposed that the electron circled the nucleus in specific orbitals like a planet around the sun. (planetary model) • The spectra produced were result of quanta emitted when the electrons “jumped” or “fell” to to lower orbitals.

23. Bohr’s Model of the Atom • Line spectra – Light from an excited element is passed through a prism. • Hydrogen • Helium • Neon • http://phys.educ.ksu.edu/vqm/html/emission.html

24. The Modern Model of the Atom • However, Bohr’s model only explained the nature of spectra of the hydrogen atom. It could not explain the spectra of an atom with more than one electron. • The charge-cloud model does not attempt to describe the path of each electron in a fixed orbit.  • Scientists now describe the possible positions of electrons in terms of probability.  • Electrons are characterized in terms of the three-dimensional shapes that their probability fields define. 

25. Evolution of Atomic Theory • Whizzy Periodic Table • Atomic Jeopardy

26. The Atom Defined • An atom is the smallest particle of an element that retains the chemical properties of that element.

27. The Atom Defined • All atoms have two regions: • The nucleus which is a very small region located near the center of the atom. • The nucleus contains protons and neutrons. • The electron cloud which is the region surrounding the nucleus that is very large compared with the size of the nucleus. • This region contains electrons.

28. Subatomic Particles • Subatomic particles are the particles that make up an atom.

29. Relative Weight If an electron weighed the same as a dime, a proton would weigh the same as a gallon of milk.

30. Reading the “Label” of an Atom • Atomic Number (Z) is equal to the number of protons in an atom. • In an atom with no charge overall (not an ion) the Atomic Number is also equal to the number of electrons.

31. Reading the “Label” of an Atom • The Mass Number (A) is equal to the number of protons and neutrons of an atom. Z = # of Protons A = # of Protons + # of Neutrons

32. Isotopes • Atoms of the same element can have different numbers of neutrons in its nucleus. • These are called isotopes. • Isotopes of an element do not differ in physical or chemical properties, only in mass.

33. Isotopes • Many elements have naturally occurring isotopes. It is the weighted average atomic mass of these isotopes that we use as the atomic mass of an element. • The unit for atomic mass is an amu (atomic mass unit). An amu = 1.66 x 10-24 g • 1 amu = 1/12th the mass of a carbon-12 atom. • There are 3 carbon isotopes. • In hyphen notation (atomic symbol-mass #) they are: C-12, C-13, and C-14 • In nuclear notation they are:

34. Isotopes • The atomic mass listed on the periodic table of the elements for each element is a weighted average atomic mass of the isotopes of that element. • So if an element has isotopes, then none of the atoms of that element actually have the atomic mass listed on the periodic table.

35. Electron Configuration – and its three rules • Electrons will occupy the lowest energy level that can receive it. This is known as the Aufbau principle. • In each energy level, electrons will fill an unoccupied orbital before an occupied level. All single electrons in orbitals will have the same spin. This is known as Hund’s rule. • No two electrons can have the same set of 4 quantum numbers (they are unique). This is known as the Pauli Exclusion Principle.

36. Electron Configuration • Principal energy level or shell: • This is the main energy level that an electron occupies. They range in numbers from 1 to 7. The lower the principal energy level number the closer an electron is to the nucleus. • Each shell has sublevels or subshells associated with it. • The number of subshells in a shell is equal to the shell number. So principle energy level 1 has 1 subshell, principle energy level 2 has 2 subshells, etc. • These subshells are designated as s, p, d, and f.

37. Electron Configuration • Each subshell has 1 or more orbital in it. • The s subshell has 1 orbital • The p subshell has 3 orbitals • The d subshell has 5 orbitals • The f subshell has 7 orbitals • Each orbital can contain 2 electrons. • These 2 electrons must spin in opposite directions.

38. Shapes of the Subshells • s-orbital p-orbital • d-orbital f-orbital