1. Atoms CHAPTER 4 1

2. Let’s Review! o Matter is… o Anything that has mass and takes up space o All matter is made of elements – substances that cannot be broken down o And elements are made of – ATOMS! o The atom is the SMALLEST unit of matter 2

3. 3 ATOMS ARE EVERYWHER E! They are in the air you breathe, the chair you’re sitting on, and the clothes you’re wearing.

4. Atomic Theory SECTION 1 4

5. DemocritusDemocritus (400 B.C.) o Ancient Greece o Theorized that matter could not be divided infinitely, you had to reach a smallest piece o Atomos: indivisible or can’t be cut (becomes atom) o Since this was ancient Greece, he had no proof and few people believed it. 5

6. John Dalton (1808) o People started to accept his idea of atoms because of his experiments o He thought: o all elements are made of atoms (indivisible and indestructible) o atoms of the same element are exactly alike o atoms of different elements are different 6

7. Law of Definite Proportions o A chemical compound is made up of the same percentage of elements. o For example, water is always made up of 2 parts hydrogen to 1 part oxygen. o Therefore we have the equation oH2OoH2O o This is true for all chemical compounds! o This law supported Dalton’s theory of atoms. 7

8. Sir J. J. Thomson (1897) o Conducted an experiment that showed atoms could be divided o Used cathode ray experiments to show negatively charged particles that came from inside atoms o These are electrons! 8

9. Cathode Rays o There are two metal plates at the end of a vacuum tube o Cathode, which has a negative charge o Anode, which has a positive charge o When voltage is applied across the plates, a glowing beam comes from the cathode and strikes the anode o Since it was a vacuum tube (a tube with all the air vacuumed out), he could see that the beam came from the negative end of the cathode 9

10. Plum Pudding o He proposed that electrons are spread throughout the atom just like blueberries are in a muffin o Plum Pudding Model 10

11. Thomson o Found negative particles could come from neutral elements o Atom is made of smaller things (+ & -), and is divisible o Successfully separated negative particles (electrons) but could not separate the positive particle 11

12. Ernest Rutherford (1911) o Fired positively charged particles at a sheet of gold foil o Most went through unaffected, some bounced away o His experiments suggested that an atom’s positive charge was concentrated at the center of the atom o This is the NUCLEUS! 12

13. Ernest Rutherford o Electrons are scattered near the outside of the atom with mostly empty space between the nucleus and the electrons o Compared to the atom, the nucleus is very small o Rutherford’s experiments led to a new model of the atom 13

14. Niels Bohr (1913) o Rutherford proposed that electrons orbited the nucleus o Similar to planets orbiting the sun o Bohr discovered the electrons are actually in energy levels (orbitals) o Their path around the nucleus is limited to their energy o Only certain electrons can be in certain levels 14

15. Modern Atomic Theory Today we know that atoms have ◦A nucleus with ◦Protons ◦Neutrons ◦Electrons that orbit the nucleus 15

16. Atomic Structure SECTION 2 16

17. Subatomic Particles o Subatomic: lower (or smaller) than an atom o Protons: positive charge o Neutrons: no charge (neutral) o Electrons: negative charge o Protons and neutrons are both in the nucleus of the atom. o Electrons are buzzing around the outside in the electron cloud 17

18. Mass and Volume o The nucleus makes up 99.99% of the mass of the atom. o However, the nucleus is 1/100,000 of the volume of an atom. ◦The volume is determined by the electron cloud. 18

19. Mass of Particles Mass of Particles o Since subatomic particles are so small they cannot be measured in grams o They are measured in atomic mass units or amu o 1 amu = 1.61x10 -24 g o 0.00000000000000000000000161! o 1 g is about the mass of a paper clip! 19

20. Mass of Particles 20 namesymbolchargelocationmass protonsp+p+ positivenucleus1 amu neutronsnono neutralnucleus1 amu electronse-e- negativeelectron shell.0006 amu

21. How do we know how many protons, neutrons, and electrons are in an atom? 21

22. Atomic Number & Atomic Mass Atomic Number: ◦number of protons found in the nucleus (therefore, it is also the number of electrons orbiting the nucleus) Atomic Mass: ◦weight of the protons and the neutrons combined. 22

23. Let’s Try! 23 Atomic # Mass Protons Neutrons Electrons Atomic # Mass Protons Neutrons Electrons Atomic # Mass Protons Neutrons Electrons Atomic # Mass Protons Neutrons Electrons 14 28 14 9 19 9 10 9 90 232 90 142 90 3734337343

24. Charges o The # of protons in an atom is unique to each element (this is how they are identified) o If an atom is stable, there will be an equal number of protons and electrons ◦This means there will be NO overall charge on the atom o The number of neutrons in an atom can vary also! 24

25. Electrons o Electrons fill orbitals around the nucleus o Each orbital can only hold so many electrons o The number of levels filled depends on how many electrons the atom has o The electrons that fill the outermost orbital are called valence electrons 25

26. Isotopes o The atomic number, or number of protons, will never change, but the number of neutrons CAN change o Isotope: Any two atoms of the same element with a different number of neutrons. o Isotopes have the SAME chemical and physical properties. o There are normally one or two stable isotopes for an element o All others are unstable (they fall apart) o radioactive decay 26

27. Representing Isotopes o Isotopes are shown in isotope notation o Since isotopes have a different atomic mass, you must write the symbol of an element AND write its atomic mass 27

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