1. History of the Atom

2. Atoms and Elements Any material that is composed of only one type of atom is called an element. An atom is the smallest particle of any element that still retains the characteristics of that element. There are 94 elements that are naturally occurring (at least in trace amounts)

3. Timeline of the Atom 460 – 370 BC: Democritus “atomos” (indivisible & indestructible) Democritus’ thoughts about atoms being indestructible were accepted for nearly 2000 years

4. 1743 – 1794: Lavoisier Father of Modern Chemistry Turned chemistry from a science of observation into a science of measurement with his carefully controlled experiments Proved that matter is neither created nor destroyed Identified and named many elements, like oxygen and hydrogen

5. 1766 – 1844: John Dalton Father of Modern Atomic Theory Studied the ratios in which elements combined in chemical reactions Determined that elements may combine in more than one proportion, forming different compounds Devised a system of rules that describe the behavior of atoms, known as Dalton’s Atomic Theory

6. Dalton’s Atomic Theory 1. All elements are composed of tiny indivisible particles called atoms. 2. Atoms of the same element are alike. The atoms of any one element are different from those of any other element. 3. Atoms of different elements can chemically combine in whole-number ratios to form compounds. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element cannot change into atoms of a different element as a result of a chemical reaction.

7. 2.1

8. 1897: JJ Thomson Discovered electrons using cathode ray tube. He showed electrons had a negative charge by deflecting the cathode ray beam (electron stream) with positive electrical charges.

9. Plum Pudding Model of the Atom Most scientists in Thomson’s day thought that the electrons were evenly distributed throughout the atom, like the plums in a plum pudding

10. 1868-1953: Robert Millikan Determined the quantity of charge carried by an electron Could then calculate the mass of the electron (very tiny – 1/1840 the mass of a hydrogen atom)

11. 1911: Ernest Rutherford Based on his work with firing alpha particles at gold foil, he proposed that the atom is mostly empty space – disproving the plum pudding model He concluded that all the positive charge and almost all the mass of an atom were concentrated in a small central region he called the nucleus. Rutherford’s model is known as The Nuclear Atom

12. Rutherford’s Gold Foil Experiment

13. Rutherford’s Model of the Atom “The Nuclear Atom”

14. 1932: James Chadwick Confirmed the existence of a subatomic particle w/ no charge – the neutron Mass nearly equal to proton (slightly larger)

15. Nuclear particles The nucleus is composed of protons and neutrons. So what are protons and neutrons made up of? A quark is an elementary particle of matter Quarks combine to form composite particles called hadrons, the most stable of which are protons and neutrons

17. Distinguishing Between Atoms Atomic Number - the number of protons in the nucleus of the atom. Each element has a unique atomic number Atomic Mass – The mass of the nucleus, which includes both protons and neutrons. Periodic tables give the average atomic mass for all the atoms of a given element. Mass Number – The total number of protons and neutrons found in a specific atom.

18. Isotopes Atomic number for all atoms of any given element is always the same. Mass number for an atom will vary depending on how many neutrons are present in that specific atom Isotopes are atoms that have the same number of protons (therefore are the same element), but different number of neutrons Example : 14 C = 6 p +, 8n o 12 C = 6 p +, 6n o

19. Atomic Number vs Mass Number

21. Isotope Practice A nitrogen atom with 8 neutrons is written: A nitrogen atom with 7 neutrons is written: A phosphorus atom with 15 neutrons is written: A phosphorus atom with 17 neutrons is written:

22. Ions An atom is electrically neutral; it has the same number of protons (positive charge) as electrons (negative charge). An ion is an atom that has a positive or negative charge because it has lost or gained an electron Sodium atom: Na (sodium) 11p +, 11e - Sodium ion:Na + (sodium ion) 11p +, 10e -

24. Ion Practice A potassium ion with 18 electrons has a charge of ? number of protons = number of electrons = overall charge = An oxygen ion with 10 electrons has a charge of ? number of protons = number of electrons = overall charge =

25. Average Atomic Mass (found on the periodic table) Found using mass spectrometer Based on the weighted average mass of all the atoms in a naturally occurring sample of the element (including all isotopes). This is why the atomic mass on a periodic table is not usually a whole number! Measured in amu’s (atomic mass unit) One amu has been defined as 1/12 the weight of a carbon atom (roughly equal to the weight of one proton or one neutron)

26. The atomic mass isn’t usually a whole number! All lithium atoms have 3 protons Some lithium atoms have 3 neutrons, and some have 4 neutrons. The atomic mass of an individual lithium atom could be 6 or 7 depending on the number of neutrons The atomic mass reported on a periodic table is the average of the masses of all existing isotopes of lithium – most of them have 4 neutrons, so the average mass is closer to 7 than 6.

27. Mass Spectrometer Used to measure atomic mass – Can identify unknown compounds by their mass – Can help identify the structure of the compound by studying the masses of the fragments the compound breaks into Works by ionizing the compound, then deflecting the various ions with a magnetic field. You can determine the mass of the compound based on how much deflection you get with a given magnetic field.

28. Mass Spectrometer

29. To calculate average atomic mass we need to know three things: 1.How many different isotopes of the element there are 2.The % of each isotope present (abundance) 3. The mass of each isotope

30. Calculating Atomic Mass Element X has 2 naturally occurring isotopes: 10 X 11 X Mass = 10.012 amuMass = 11.009 amu abundance= 19.91%abundance = 80.09% _______________ For 10 X: 10.012 x 0.1991 = 1.993 amu For 11 X: 11.009 x 0.8009 = 8.817 amu (add) Mass of element X =10.810 amu

31. Calculating Atomic Mass 35 Cl75% abundance 37 Cl 25% abundance (Will the atomic mass be closer to 35 or 37?)

32. Calculating Atomic Mass 63 Cu69.2% abundance 65 Cu30.8% abundance