1. Atoms, Molecules and Ions Chapter 2

2. Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.

3. Dalton's Postulates Each element is composed of extremely small particles called atoms. (pg 38)

4. Dalton's Postulates All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

5. Dalton's Postulates Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. Law of Conservation of Mass

6. Dalton’s Postulates Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

7. J.J. Thomson, measured charge/mass of e - (1906 Nobel Prize in Physics) Figure 2.4

8. Fig. 2.3(b)

9. Fig. 2.3(c)

10. The Atom, circa 1900 The prevailing theory was that of the “plum pudding” model, put forward by Thompson. It featured a positive sphere of matter with negative electrons imbedded in it.

11. Figure 2.5 Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other. Millikan’s Oil-drop Experiment Measured mass of e - (1923 Nobel Prize in Physics)

12. (Uranium compound) Figure 2.8

13. 1.Atom’s positive charge is concentrated in the nucleus 2.Proton (p) has opposite (+) charge of electron (-)  particle velocity ~ 1.4 x 10 7 m/s (~5% speed of light) (1908 Nobel Prize in Chemistry) Figure 2.10

14. The Nuclear Atom Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. Most of the volume of the atom is empty space.

15. atomic radius ~ 100 pm = 1 x 10 -10 m nuclear radius ~ 5 x 10 -3 pm = 5 x 10 -15 m Rutherford’s Model of the Atom

16. Subatomic Particles (Table 2.1) ParticleChargeRelative Mass (amu) electron (e-)≈ 0 proton (p)+11 neutron (n)01

17. Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes - atoms of the same element (X) with different numbers of neutrons in their nuclei X A Z H 1 1 H (D) 2 1 H (T) 3 1 U 235 92 U 238 92 Mass Number Atomic Number Element Symbol

18. How many protons, neutrons, and electrons are in C 14 6 ? How many protons, neutrons, and electrons are in C 11 6 ? 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons Do You Understand Isotopes?

19. Average Mass Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. Average mass is calculated from the isotopes of an element weighted by their relative abundances. Average mass is calculated from the isotopes of an element weighted by their relative abundances.

20. The Three Isotopes of Hydrogen 99.9885%0.0115%Trace%

21. H 1 1.00794

22. Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. Fig 2.15

23. Period Group Alkali Metal Noble Gas Halogen Alkali Earth Metal Figure 2.16