1. S.MORRIS 2006

2. Where did it all begin? The word “atom” comes from the Greek word “atomos” which means indivisible. The idea that all matter is made up of atoms was first proposed by the Greek philosopher Democritus in the 5th century B.C.

3. HISTORY OF THE ATOM 460 BC Democritus develops the idea of atoms he pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called ATOMA (greek for indivisible)

4. HISTORY OF THE ATOM 1803 John Dalton suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS

5. Dalton’s Atomic Theory (1808) 1. Elements are composed of extremely small particles called atoms. 3. Atoms of different elements can combine in simple whole number ratios to form compounds. 4. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions. 2.1 2. All atoms of a given element are identical. The atoms of one element are different from those of any other element

6. 8 X 2 Y 16 X8 Y + 2.1

7. Law of Definite Proportions The ratio of mass of elements in a compound is always the same Every Water molecule will contain 16g of oxygen and 2 g of hydrogen

8. 2 2.1

9. HISTORY OF THE ATOM 1898 Joseph John Thompson found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON

10. HISTORY OF THE ATOM Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge 1890 like plums surrounded by pudding. PLUM PUDDING MODEL

11. HISTORY OF THE ATOM 1910 Ernest Rutherford oversaw Geiger and Marsden carrying out his famous experiment. they fired Helium nuclei at a piece of gold foil which was only a few atoms thick. they found that although most of them passed through. About 1 in 10,000 hit 1910

12. 1.atoms positive charge is concentrated in the nucleus 2.proton (p) has opposite (+) charge of electron 3.mass of p is 1840 x mass of e - (1.67 x 10 -24 g)  particle velocity ~ 1.4 x 10 7 m/s (~5% speed of light) (1908 Nobel Prize in Chemistry) 2.2

13. atomic radius ~ 100 pm = 1 x 10 -10 m nuclear radius ~ 5 x 10 -3 pm = 5 x 10 -15 m Rutherford’s Model of the Atom 2.2

14. HISTORY OF THE ATOM gold foil helium nuclei They found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back. helium nuclei

15. HISTORY OF THE ATOM Rutherford’s new evidence allowed him to propose a more detailed model with a central nucleus. He suggested that the positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction However, this was not the end of the story.

16. HISTORY OF THE ATOM 1913 Niels Bohr studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons.

17. Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4  + 9 Be 1 n + 12 C + energy neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x 10 -24 g 2.2

18. Subatomic Particles mass p = mass n = 1840 x mass e - 2.2

19. Bohr’s Atom electrons in orbits nucleus

20. Atoms # of protons & neutrons The basic unit of Matter The smallest particle of an element that retains the properties of that element.

21. HELIUM ATOM + N N + - - proton electron neutron Shell What do these particles consist of?

22. ATOMIC STRUCTURE All About Atoms All About Atoms Particle Proton Neutron electron Charge + charge - charge No charge 1amu 1/1836 Mass

23. ATOMIC STRUCTURE Represents the number of protons in an atom Never changes P+ equal to the number of e- the number of protons and neutrons in an atom Neutrons equal mass # - atomic # He 2 4 Mass Number Atomic number number of electrons = number of protons

24. Ions Charged particles due to the loss or gain of electrons 2) Atoms that are called cations lose e- thus becoming positive 1) Atoms that are called anions gain e- thus become negative Na 11 protons 11 electrons Na + 11 protons 10 electrons Cl 17 protons 17 electrons Cl - 17 protons 18 electrons 2.5

25. Isotopes Atoms with the same atomic number but different mass number Atoms having the same number of protons but different numbers of neutrons Average Atomic Mass of an element is the weighted average of an element’s naturally occurring isotopes

26. 2.3

27. How many protons, neutrons, and electrons are in C 14 6 ? How many protons, neutrons, and electrons are in C 11 6 ? 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons Do You Understand Isotopes? Isotope Maker 2.3

28. Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons # OF NEUTRONS = mass number – atomic number X A Z H 1 1 H (D) 2 1 H (T) 3 1 U 235 92 U 238 92 Mass Number Atomic Number Element Symbol 2.3

29. Bohr Model of the Atom Electrons travel around the nucleus in one of several orbits/shells/ nrg levels Principal energy level is designated by a quantum number (n) Quantum number is the same as period on the periodic table

30. Period Group Alkali Metal Noble Gas Halogen Alkali Earth Metal 2.4

31. ATOMIC STRUCTURE Electrons are arranged in Energy Levels or Shells around the nucleus of an atom. first shell maximum of 2 electrons second shell maximum of 8 electrons third shell max of 18 electrons fourth shell max of 32 electrons

32. ATOMIC STRUCTURE There are two ways to represent the atomic structure of an element or compound; 1. Electronic Configuration 2. Electron dot diagrams

33. ELECTRONIC CONFIGURATION With electronic configuration elements are represented numerically by the number of electrons in their shells and number of shells. For example; N Nitrogen 14 7 2 in 1 st shell 5 in 2 nd shell configuration = 2, 5 2 + 5 = 7

34. ELECTRONIC CONFIGURATION Write the electronic configuration for the following elements; Ca O Cl Si Na 20 40 11 23 8 17 16 35 14 28 B 11 5 a)b)c) d)e)f) 2,8,8,22,8,1 2,8,72,8,42,3 2,6

35. Valence Valence shell is outermost occupied energy level and is the same as the period number in the periodic table Valence electrons are the electrons in the outer energy level of an atom

36. DOT & CROSS DIAGRAMS With Dot & Cross diagrams elements and compounds are represented by Dots or Crosses to show electrons, and circles to show the shells. For example; Nitrogen N XX X X XX X N 7 14

37. DOT & CROSS DIAGRAMS Draw the Dot & Cross diagrams for the following elements; OCl 817 16 35 a)b) O X X X X X X X X Cl X X X XX X X X X X X X X X X X X X

38. SUMMARY 1. The Atomic Number of an atom = number of protons in the nucleus. 2. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. 3. The number of Protons = Number of Electrons. 4. Electrons orbit the nucleus in shells. 5. Each shell can only carry a set number of electrons.