Chapter 3

1. Chapter 3 Atoms: The Building Blocks Of Matter

2. Section 3-1 The Atom: From Philosophical Idea To Scientific Theory

3. Two Theories on the Nature of Matter • It is continuous and can be infinitely broken down into smaller and smaller parts. • It has some fundamental particle that can be reached that cannot be broken down any further. • The idea of the fundamental particle is the one accepted today.

4. The study of matter lead to two basic laws of chemistry • The Law of Conservation of Matter – During ordinary chemical and physical reactions, matter is neither created nor destroyed, it is conserved.

5. The study of matter lead to two basic laws of chemistry • The Law of Definite Proportions – A chemical compound contains the same elements in exactly the same proportions by mass, regardless of the size of the sample or the source of the compound.

6. John Dalton – (1766-1844) • English Schoolteacher • Proposed an explanation for the Law of Conservation of Mass and the Law of Definite Proportions in 1808. • This was called Dalton’s Atomic Theory. It has five statements.

7. Dalton’s Atomic Theory • All matter is composed of submicroscopic, indivisible particles called atoms. • Atoms of the same element are the same, atoms of different elements are different.

8. Dalton’s Atomic Theory • Atoms cannot be broken down, created, destroyed, or changed into other atoms. • Atoms of different elements combine in whole number ratios to form compounds.

9. Dalton’s Atomic Theory • In chemical reactions, atoms are combined, separated, or rearranged.

10. Dalton’s Atomic Theory is not Completely Correct • Atoms are divisible. • Atoms of the same element are not always the same. • Atoms are not submicroscopic. • Atoms can be changed into other atoms.

11. Section 3-2 The Structure of the Atom

12. Structure of the Atom • An atom is the smallest particle of an element that retains the properties of that element. • It contains two regions, the nucleus and the electron cloud.

13. The Nucleus Small, Dense region in the center of an atom. Holds at least one proton, most hold at least one neutron. The nucleus is positively charged. The Electron Cloud Region around the nucleus that contains electrons. The Regions of the Atom

14. The Three Subatomic Particles • Electrons – Negatively charged, found outside of the nucleus. Smallest subatomic particle. • Protons – Positively charged, found in the nucleus. Middle sized subatomic particle. • Neutron – No charge, found in the nucleus. Largest subatomic particle.

15. Discovery of the Electron • In the late 1800’s many experiments were done by passing current through gases at low pressure. • These were carried out in sealed glass tubes called cathode ray tubes.

16. Discovery of the Electron • When a current was passed through the tube, the anodes (+) glowed. • The theory was that the particles that made up the cathode ray were negatively charged, and moved from cathode to anode.

17. Joseph John Thompson (1856 – 1940) • Conducted experiments to support this idea in 1897. • Called these particles corpuscles. • Determined it had a very large negative charge with a very small mass. • Conclusion was that all atoms contained electrons.

18. The Thompson Model of the Atom • Thompson’s model of the atom is different from Dalton’s because it shows negative electrons stuck in positive ball of matter. • There are two problems. First, it does not show how easy it is to remove electrons, and second, it does not show numbers of electrons.

19. Assumptions from Thompson’s Experiments • Atoms must be divisible. This proves part of Dalton’s theory wrong. • Since atoms are neutral, there must be a positively charged particle to counteract the negatively charged electron. • Because electrons are so small, there must be another particle that gives the atom its mass.

20. Lord Ernest Rutherford (1871 – 1937) • Thompson’s graduate student. • He, Ernest Marsden and Hans Geiger worked together firing alpha particles at a sheet of gold foil. • Observed that most of the particles passed through the foil, but some were deflected back at them.

21. Discovery of the Nucleus • To the scientists, this was like firing a gun at a piece of tissue paper and having it bounce off back at them. • Rutherford theorized that the particles were bouncing off a small, dense, positively charged region in the atom. • He called this the nucleus.

22. The Rutherford Model of the Atom • Rutherford developed a new atomic model in 1913. • In it, the electrons move in space around the positively charged nucleus. • The problem with his model is that the nature of electrons is not accurately described by this model.

23. The Atomic Nucleus • The nucleus contains protons and neutrons. • It is the number of protons that determines the identity of an element. • Atoms with more than one proton should not exist. • Nuclear forces hold the nucleus together. These are strong attractions between protons and protons, neutrons and neutrons, and protons and neutrons.

24. Atoms are very small and Numerous. • Neon for example is a gas that makes up only 0.002% of our air. • In each breath you take, you breath in 5 * 1017 atoms of neon. • If the nucleus of an atom were increased in size to 1cm3, it would have a mass of 2 * 108 tons.

25. Section 3-3 Counting Atoms

26. There are three numbers that identify an atom • Atomic Number • Mass Number • Atomic Mass

27. Atomic Number • The number of protons in the nucleus. • This is the number that identifies an element. • This is the whole number on the periodic table.

28. Mass Number • Most elements have isotopes. • Isotopes – Atoms of the same elements with different numbers of neutrons. • Mass Number – The number of protons plus the number of neutrons. • There are two ways to write the isotope of an element. • Write the name, then a hyphen, then the mass number – Carbon – 14 • Write the symbol, mass number as a superscript, atomic number as a subscript. • An isotope is also called a nuclide.

29. Atomic Mass • Do not use grams to measure mass of an atom, instead we use atomic mass units (amu). • Average atomic mass is the weighted average of the masses of all of the isotopes of an element. • This is the decimal number on the periodic table.

30. There are three ways to compare numbers of atoms. • Moles • Avogadro’s Number • Molar Mass

31. The Mole • The amount of substance that contains as many particles as there are in 12 grams of carbon – 12.

32. Avogadro’s Number • The number of particles in one mole of a pure substance. • Named after Italian scientist Amadeo Avogadro (1776 – 1856). • 6.02 * 1023 particles.

33. Molar Mass • The mass of one mole of a pure substance. • Expressed in grams. • It is equal to the atomic mass (decimal number on the periodic table) in grams.