Atomic Structure

1. Atomic Structure

2. Atom • Definition – smallest particle of an element that retains the chemical identity of that element • Term proposed by Democritus in 450 BC • Democritus proposed that all matter is composed of tiny, indivisible particles • Rejected by Aristotle because Democritus could not explain how particles “stuck together”

3. Dalton’s Atomic Theory of Matter- 1803 • Based on observations made by scientists during the 1700’s, John Dalton proposed a theory that could explain the properties of matter • Four postulates: • Each element is made of small particles called atoms • All atoms of a given element are identical, but different from those of any other element • Atoms are neither created nor destroyed in a chemical reaction • A given compound always has the same relative numbers and kinds of atoms

4. Are the four postulates all true? • What do we know about atoms that means we might need to re-state one of the postulates?

5. Postulate Number 2 – not all atoms of the same element are the same • You can have different Isotopes of the same element Example: Iodine-125 and Iodine-127 are both iodine, but are different from each other. Iodine-125 is radioactive and Iodine-127 is not.

6. History of Discovering about Atomic Structure • Dalton thought atoms were hard and round like tiny marbles • Early 1800’s – Franklin- two kinds of electric charges (positive and negative)

7. Electrons were First Particles Found • An English scientist discovered electrons and named them. An Amercian scientist found their mass. Mass of an electron to be 9.11 x 10-28 grams

8. Protons were next particles found • Atoms are electrically neutral so the negative electrons must be balanced by an equal number of positive particles • An English scientist found these particles and called them protons. • He also discovered these positively charged particles were all concentrated in the center of the atom – called it the nucleus

9. Modern Atomic Theory • Atoms are composed of protons, neutrons, and electrons • At first, scientists visualized electrons orbiting the nucleus much like planets orbiting the sun

10. Last particle found was the neutron • Neutrons have a mass approximately equal to that of a proton • Neutrons have no charge • Neutrons help “glue” the nucleus together

11. Fundamental Subatomic Particles ParticleLocationCharge Mass (amu) Proton Nucleus +1 1 Neutron Nucleus 0 1 Electron Outside -1 0 Nucleus

12. Moseley’s Discovery • Found that atoms of each element contain a unique positive charge in their nucleus • Conclusion: atom’s identity comes from the number of protons in the nucleus Called “Atomic Number” • Neutral atoms – positive charges must equal negative charges Conclusion – for neutral atoms, no. of electrons equals no. of protons

13. Information in Periodic Table 7 Atomic Number (sometime not in center) N Chemical Symbol Nitrogen 14.0067 Atomic Mass

14. Mass Number • Hydrogen -1 The 1 after the hydrogen is the mass number. Used to help keep track of what isotope you have. • Definition Sum of the neutrons and protons Always a whole number • Another way to represent Mass Number1 2 3 H H H 1 1 1 Atomic Number

15. Mass of an Atom • Atomic Mass Unit – approximately the weight of one neutron or proton • Used because the weight in grams is very small 1 a.m.u. = 1.66 x 10-24 g.

16. Element’s atomic mass • Weighted average of all of the isotopes of an element 7 N Nitrogen 14.0067 Atomic Mass Nitrogen contains mostly Nitrogen-14 with a small amount of Nitrogen-15

17. Practice • Find sodium (Na) on the periodic table I gave you • Find the atomic number. What does that number mean? • Find the atomic mass. What does that number mean?

18. How many electrons does one sodium atom have? • I have one atom of sodium represented by 46V 23 How many protons does it have? How many neutrons does it have?