1. Structure of the Atom: The Developing Story SNC1D/E

2. Early Views 2500 years ago, Greek philosopher Democritus figured that matter was made up of small particles This was rejected by Aristotle and others Aristotle figured all matter consisted of either earth, air, water or fire Took almost 2000 years to return to the theory that matter was made of small particles

3. Dalton’s Atom 1808 Indivisible Sphere

4. Dalton ’ s Atomic Theory (1808): 1. All matter is made up of tiny, indivisible particles called atoms. 1. All atoms of an element are identical. 1. Atoms of different elements are different. 1. Atoms are rearranged to form new substances in chemical reactions, but they are never created or destroyed. Which are still considered true today? For those postulate which have changed, how are they different?

5. J.J. Thompson 1897

6. Thomson’s Model

7. Positively charged sphere, with negative electrons distributed EVENLY throughout the atom.

8. Ernest Rutherford 1920 Note that this is not to scale (it would not fit on the slide – why?)

9. Rutherford ’ s Gold Foil Experiment http://www1.teachertube.com/viewVideo.php?title=Rutherford__s_experiment&video_id=8329  Rutherford aimed a beam of positively charged particles (alpha particles) at a sheet of gold foil.  The gold foil was surrounded by a fluorescent screen, which would glow when the alpha particles hit it.  If Thomson’s model were true, the positive particles would travel either straight through the evenly distributed charges in the gold atoms, or straight through the spaces between the atoms, and hit the screen on the opposite side.  Most particles did travel straight through, but some deflected, and others even bounces straight back!

10. Ernest Rutherford (1920) All of the positive charge in an atom is located in the centre of the atom, called the nucleus. The positively charged particles in the nucleus are called protons. The nucleus contains most of the atom’s mass, but takes up very little space. The nucleus is surrounded mostly by empty space which is occupied by electrons.

11. Neils Böhr Planetary Model

12. Neils Böhr Passed a current through a vacuum tube containing hydrogen, which glowed with four SPECIFIC colours Proposed electrons could only exist at SPECIFIC energy levels Explanation: Electrons were “excited” by electrical energy and would jump to higher energy levels. As electrons return back down to their original level, they would release energy as light. Depending on how many energy levels the electron came down, a specific energy of light would be emitted.

13. Niels Böhr (1922) The electrons orbit the nucleus like planets orbit the Sun. Each electron orbit has a defined amount of energy. Electrons can only exist in specific energy levels. Electrons cannot be between the orbits, but they can “jump” from one orbit to another. When atoms release energy it is seen as light. Each orbit can hold a maximum number of electrons.

14. Neils Böhr Planetary Model

15. James Chadwick (1932) Confirmed that the nucleus contained neutral particles, with about the same mass as protons

16. Formation of the Atomic Theory

17. The Modern Atom Atoms are the smallest particles of an element. Atoms are made up of subatomic particles (proton, electrons, neutrons). Protons and neutrons are found in the nucleus. Electrons occupy space around the nucleus. Protons and neutrons make up the majority of the atoms mass. Example A proton has a mass of about 1 atomic mass unit (u), which is equal to 1.66×10 -10 g.

18. The Atom The atom is made up of three subatomic particles:  Protons (p + )  Electrons (e - )  Neutrons (n o )

19. The Atom ParticleLocationChargeMass Protonnucleus+11 Electrons outer shells 0 Neutronsnucleus01

20. The Atom Atomic number  the number of protons in the nucleus of an atom  The atomic number identifies the atom as a particular element Mass Number/Atomic Mass  The sum of the number of protons and the number of neutrons in an atom

21. On the Periodic Table vs Standard Atomic Notation

22. The Atom Calculating the Number of Subatomic Particles: Atomic number = number of protons Mass number = number of protons + number of neutrons Number of neutrons = mass number – atomic number Number of protons = number of electrons (neutral atom)

23. Example Atomic # = Mass # = Standard Atomic notation = “ pen ” # of protons = # of electrons = # of neutrons =

24. Example Atomic # = Mass # = Standard Atomic notation = “pen” # of protons = # of electrons = # of neutrons =

25. Elements to Know  You are responsible for knowing the name and symbols for the first 20 elements and the following extras:  Chromium (Cr)  Iron (Fe)  Cobalt (Co)  Nickel (Ni)  Copper (Cu)  Zinc (Zn)  Bromine (Br)  Silver (Ag)  Tin (Sn)  Iodine (I)  Gold (Au)  Mercury (Hg)  Lead (Pb)  Barium (Ba)