1. Atoms: The Building Blocks of Matter

2. PART 1

3.  Democritus [400 B.C]  Greek philosopher  Hypothesized: Nature has a basic indivisible particle of which everything is made of Called this particle an atom Greek “atomos” = indivisible

4.  Law of Conservation of Mass  Mass is neither created nor destroyed during ordinary chemical reactions or physical changes  Law of Definite Proportions  A chemical compound contains the same elements in exactly the same proportions by mass regardless of size of sample or source of compound i.e. Every sample of table salt is made of 39.34% Na and 60.66% Cl i.e. H 2 O always has 2 atoms of H and 1 atom of O

5.  Law of Multiple Proportions  If two or more different compounds are composed of the same two 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 i.e. CO and CO 2 CO = 1.00g of C and 1.33 g of O CO 2 = 1.00 g of C and 2.66 g of O The ratio of the second element is 2.66 to 1.33 or 2 to 1

6.  John Dalton [1808]  English schoolteacher – liked nature and weather  Developed: Dalton’s Atomic Theory

7. 1. All matter is composed of extremely small particles called atoms 2. Atoms of a given element are identical in size, mass and other properties and are different from atoms of other elements 3. Atoms cannot be subdivided, created, or destroyed 4. Atoms of different elements combine in simple whole number ratios to form chemical compounds 5. In chemical reactions, atoms are combined, separated or rearranged

8.  Atoms can be split into even smaller particles (nuclear chemistry) and aren’t indivisible  i.e. nucleus, protons, electrons  A given element can have different masses  i.e. isotopes

9.  Today’s definition of the atom  Atom = Smallest particle of an element that retains the chemical properties of that element Two regions Nucleus  Very dense, small center of the atoms  Protons and neutrons Electron Cloud  Region occupied by electrons  Subatomic particles  Protons, neutrons, electrons

10.  J.J. Thomson [1897]  Discovered: The 1 st subatomic particle: the negatively charged electron  Used a Cathode Ray Experiment Cathode Ray Tube – Electric current passed through a metal disk to another metal disk in a gas at low pressure (vacuum sealed tube) i. e. neon signs and ‘old- fashioned’ television sets

11.  When a current passed through the cathode ray tube, the surface of the tube opposite the cathode glowed  Glow was hypothesized to be stream of particles called a cathode ray  Ray affected by magnetic fields Attracted to positive charge Deflected from negative charge http://www.youtube.com/watch?v =7YHwMWcxeX8&NR=1

12.  Thomson measured the ratio of the charge of the particles to their mass  Same ratio no matter what metal or gas was used  Named this particle an electron  http://www.youtube.com/watch?v=IdTxGJjA4Jw&feature=related http://www.youtube.com/watch?v=IdTxGJjA4Jw&feature=related

13.  Atoms are electrically neutral  Must have positive charges to balance the negatively charged electrons  Electrons have a lot less mass than atoms  Other particles must account for their mass  Plum Pudding Model  positively charged sphere with electrons dispersed through it

14.  Robert Millikan [1909]  Discovered: The measurement of an electron charge  Oil Drop Experiment Measured the difference in velocity of oil droplets Charged droplets (ionizing radiation) vs. uncharged  http://www.youtube.com /watch?v=XMfYHag7Liw&fe ature=related http://www.youtube.com /watch?v=XMfYHag7Liw&fe ature=related

15.  Ernest Rutherford (with Hans Geiger and Ernest Marsden) [1911]  Discovered: A new atomic model  Gold Foil Experiment Bombarded thin piece of gold foil with alpha particles Expected alpha particles to pass through with minimal deflection Surprised when 1 in 8000 deflected back to source It was “as if you had fired a 15 inch artillery shell at a piece of tissue paper and it came back and hit you” http://w ww.yout ube.com /watch?v =wzALbz Tdnc8&f eature=r elated

16.  Discovered the nucleus is a small densely packed volume of positive charge  Size comparison Nucleus = marble Whole Atom = football field  At this point in history, we were not sure where the electrons were – stay tuned for more in Chapter 4

17. PART 2

18.  2 types of particles  Protons positively charged = +1 made up of quarks  Neutrons neutral = 0 charge Made up of quarks  Mass in the nucleus Protons = 1.673 x 10 -27 Neutrons = 1.675 x 10 -27 To simplify, both have mass of 1 amu (atomic mass unit)

19.  Strong Nuclear Forces  Two protons extremely close = strong attraction  Two neutrons extremely close = strong attraction  Neutrons and Protons extremely close = strong attraction  Strong nuclear forces overcome the repulsion of like positive charges to keep the nucleus together!!!

20.  In the Electron Cloud  A cloud of negative charge outside of the nucleus  More on this later........  Electrons = Negatively charged particles with almost no mass (9.109 x 10 -31 )

21.  Atomic Number  Equal to the number of protons and specific to each type of element  Identifies the element # of protons is what give that element its characteristic properties Elements with different protons are NOT THE SAME ELEMENT!!!

22.  Neutral atoms  total positive charge equals the total negative charge # protons (+1 each) = # electrons (-1 each)

23.  Atoms of the same element (i.e. same # of protons) that have differing number of neutrons  Isotopes of the same element have different masses.

24.  Mass number = #protons + # neutrons  Average Atomic Mass  Every element has isotopes  The periodic table takes into account all naturally occurring isotopes of an element and averages them ElementAtomic Number # of Protons # of Neutrons Mass Number Carbon66 816 Nitrogen15

25.  Atoms with a charge  Negative – more electrons than protons  Positive – more protons than electrons  Charge = #protons - # electrons  Magnesium atom with 12 protons and 10 electrons has a charge of +2