2. How long have people been interested in understanding matter and its structure? A.Thousands of years B.Hundreds of years C.A few years D.Never

3. History of Atomic Structure

4. Ancient Philosophy Who: Aristotle, Democritus When: More than 2000 years ago Where: Greece What: Aristotle believed in 4 elements: Earth, Air, Fire, and Water. Democritus believed that matter was made of small particles he named “atoms”. Why: Aristotle and Democritus used observation and inferrence to explain the existence of everything.

5. Aristotle Democritus

6. Aristotle Early Greek Theories 400 B.C. - Democritus thought matter could not be divided indefinitely. 350 B.C - Aristotle modified an earlier theory that matter was made of four “elements”: earth, fire, water, air. Democritus Aristotle was wrong. However, his theory persisted for 2000 years. fire air water earth

7. Alchemists Who: European Scientists When: 800 – 900 years ago Where: Europe What: Their work developed into what is now modern chemistry. Why: Trying to change ordinary materials into gold.

8. Alchemic Symbols

9. Particle Theory Who: John Dalton When: 1808 Where: England What: Described atoms as tiny particles that could not be divided. Thought each element was made of its own kind of atom. Why: Building on the ideas of Democritus in ancient Greece.

10. John Dalton

11. 1800 -Dalton proposed a modern atomic model based on experimentation not on pure reason. All matter is made of atoms. Atoms of an element are identical. Each element has different atoms. Atoms of different elements combine in constant ratios to form compounds. Atoms are rearranged in reactions. His ideas account for the Law of Conservation of Mass

12. The Law of Conservation of Mass Atoms are neither created nor destroyed The Law of Constant Proportions Atoms in compounds are in fixed ratios…CO2 is a different ration than CO.

13. Discovery of Electrons Who: J. J. Thomson When: 1897 Where: England What: Thompson discovered that electrons were smaller particles of an atom and were negatively charged. Why: Thompson knew atoms were neutrally charged, but couldn’t find the positive particle.

14. J. J. Thomson

15. J.J. Thomson’s Theory Discovered a negatively charged particle in the atom, that he named the electron. “Plum Pudding Model” of the Atom. Positive and very small negative charges are distributed evenly throughout the atom.

16. Discovery of the Nucleus Who: Ernest Rutherford When: 1911 Where: England What: Conducted an experiment to isolate the positive particles in an atom. Decided that the atoms were mostly empty space, but had a dense central core. Why: He knew that atoms had positive and negative particles, but could not decide how they were arranged.

17. Ernest Rutherford

18. Most particles passed through. So, atoms are mostly empty. Some positive  -particles deflected or bounced back! Thus, a “nucleus” is positive & holds most of an atom’s mass. Radioactive substance path of invisible  -particles Rutherford shot alpha (  ) particles at gold foil. Lead block Zinc sulfide screenThin gold foil

19. Lord Rutherford’s Theory Gold Foil Experiment: –Fired alpha particles at a piece of gold foil. –Most particles when through the foil, but about 1 in 10,000 got deflected. The new evidence allowed him to formulate a new atomic model with a central nucleus. He suggested that the dense positive charge was in the center and the negatively charged electrons were being held in place by attractive forces. –Remember, opposites attract!

20. Protons and Neutrons Since atoms are electrically neutral, scientists reasoned that there must be positive charges to offset the negative charges. These positive particles were named protons It was later discovered that there are additional particles in the nucleus with no charge These neutral particles are neutrons

21. The Bohr Model Who: Niels Bohr When: 1913 Where: England What: Proposed that electrons traveled in fixed paths around the nucleus. Scientists still use the Bohr model to show the number of electrons in each orbit around the nucleus. Why: Bohr was trying to show why the negative electrons were not sucked into the nucleus of the atom.

22. Niels Bohr

23. Bohr’s model There are 2 types of spectra: continuous spectra & line spectra. It’s when electrons fall back down that they release a photon. These jumps down from “shell” to “shell” account for the line spectra seen in gas discharge tubes (through spectroscopes). Electrons orbit the nucleus in “shells” Electrons can be bumped up to a higher shell if hit by an electron or a photon of light.

24. Electron Cloud Model Electrons travel around the nucleus in random orbits. Scientists cannot predict where they will be at any given moment. Electrons travel so fast, they appear to form a “cloud” around the nucleus.

25. The Atom All matter is composed of tiny particles called Atoms. All atoms are composed of smaller particles: –Protons- atomic particles with a positive charge. –Neutrons- atomic particles with no charge. –Electrons- atomic particles with a negative charge

26. Atomic Number The atomic number- depends on the number of protons in a type of element. –In an electrically neutral atom, the number or electrons is the same at the number of protons. Number of protons = number of electrons Atomic Number Atomic Mass

27. Isotopes Atoms of the same element that have different numbers of neutrons and have different mass numbers. Isotopes can be identified by using a symbol that indicated both the element and its mass number. C-12 represents a carbon atom with a mass number of 12. C-14 is an isotope of carbon.

28. Atomic Mass Atomic mass is based on the number of protons and neutrons in an atom of an element. A certain element has an atomic mass of 16 and an atomic number of 8. The atomic number equals the number of protons in the element’s atoms. To find the number of neutrons, subtract the atomic number (8) from the atomic mass (16). 16 - 8 8 neutrons

29. Atomic Mass Atomic mass of an element is an average of all the isotopes in a sample of the element. Carbon has an atomic mass of 12.01 amu (atomic mass unit) There are two major isotopes of carbon, C-12 (98.89%) and C-13 (1.108 %).

30. Location of Electrons Electrons are found in the space of the atom around the nucleus. They are found in orbitals, the most probable location on an electron Orbitals form a series of energy levels in which electrons may be found. Each electron in an atom has its own distinct amount of energy that corresponds to the energy level.

31. Ground States When electrons occupy the lowest available orbital, the atom is said to be in the ground state. Electrons fill the available spaces from the lowest energy level to higher levels until all the electrons are accounted for. Na – 2-8-1

32. Excited State When electrons are subjected to stimuli such as heat, light, or electricity, an electron may absorb energy and temporarily move to a higher energy level. This state is the unstable condition and an excited electron quickly returns to a lower available level, emitting the same amount of energy it absorbed.

33. Excited State The energy emitted may be in the form of infrared, ultraviolet, or visible light. Light given off from a neon sign is caused by excited electrons returning to ground state

34. Electromagnetic Spectrum Light appears as one color to our eyes, however it is actually composed of many different wavelengths, each of which is seen as a different line when viewed through an instrument called a spectroscope. Unlike a continuous spectrum, visible light produced by electrons is confined to narrow lines of color called bright line spectra. Each atom has its own distinct pattern, used to identify elements.

35. Electron Arrangement The arrangement of electrons determines chemical properties These properties are based on the number of electrons in the outer energy level. These outer electrons are called valence electrons Electron arrangement 2-8-18-32

36. Quantum Numbers Quantum Theory was developed to explain the chemical behavior of atoms Quantum numbers describe the major energy levels of electrons If an electron has a principal quantum number of 2, it is in the second energy level from the nucleus

37. Quantum Numbers Each energy level has one or more sublevel associated with it. Each energy level contains as many sublevels as the number of the level. Energy level 3 has 3 sublevels.

38. Sublevels The first sublevel of any energy level is designated the s sublevel. The second is p, third is d, and fourth is f (if present) 3s describes the first sublevel of the third energy level

39. Orbitals and Orientation The third quantum number relates to the orbitals in the sublevel and their orientations. S orbitals have a spherical shape P orbitals have a dumbell shape, and each p sublevel contains three orbitals (px, py, pz)

40. Spin of the Electron The fourth quantum number relates to the spin of an electron This number indicated that each orbital can contain two electrons spinning in opposite directions.

41. Electron Configuration Quantum numbers describe the distribution of the electrons in an atom when you remember that electrons will occupy the lowest sublevel possible. The electron configuration identifies the energy level of each electron and its sublevel

42. Writing Electron Configuration 1)Each added electron is placed into the sublevel of the lowest possible energy 2)No more than 2 electrons can be placed in any orbital 3)A single electron must be placed into each orbital of a given sublevel before pairing takes place (Hund’s Rule) 4)The outermost principal energy level can only contain electrons in s and p orbitals