Do Atoms exist?

Law of Conservation of Mass The total mass of substances present at the end of a chemical process is the same as the total mass of substances present before the process took place. What are the implications of this Law?

Combust 11.5 g ethanol uses 24.0 g of oxygen = 35.5 g reactants Collect 22.0 g CO2 and 13.5 g H2O = 35.5 g of products 3.6

Law of Constant Composition Joseph Proust (1754–1826) The elemental composition as determined by mass measurements of a pure substance does not vary. Usually reported as a mass percent of each element.

Percent composition of an element in a compound = mass of element molar mass of compound x 100% When we analyze pure ethanol, the percent by mass is the following always: %C = 52.14% %H = 13.13% %O = 34.73% 52.14% + 13.13% + 34.73% = 100.0% 3.5

Multiple Proportions Some atoms (mostly nonmetals) can combine in different mass ratios with each other. In the event this happens, the mass ratios are whole numbers of each other. Link to Movie

Atomic Theory of Matter In the early 19th century, John Dalton considered what these known Laws could mean. His thoughts are now known as Dalton’s atomic theory. Figure 2.1 John Dalton (1766-1844)

Dalton’s Postulates Each element is composed of extremely small particles called atoms.

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.

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.

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. So what happened next?

The Electron Figure 2.4 Thompson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/g.

Millikan Oil Drop Experiment 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. Figure 2.5

Millikan Oil Drop Experiment Robert Millikan (University of Chicago) determined the charge on the electron in 1909. Figure 2.5 Link to movie

Radioactivity: The spontaneous emission of particles and/or high frequency waves by an atom. First observed by Henri Becquerel. Also studied by Marie and Pierre Curie.

Radioactivity Three types of radiation were discovered by Ernest Rutherford:  particles (mass about 4 times that of hydrogen)  particles (mass +- 1/2000 that of hydrogen)  rays (a high frequency wave) Link to video Figure 2.8

The Atom, circa 1900: “Plum pudding” model, put forward by Thompson. Positive sphere of matter with negative electrons imbedded in it. Figure 2.9

Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles. Figure 2.10 Link to Video

The Nuclear Atom Since some particles were deflected at large angles, Thompson’s model could not be correct. Figure 2.11

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. Figure 2.12

Other Subatomic Particles Protons were discovered by Rutherford in 1919. Neutrons were discovered by James Chadwick in 1932.

Subatomic Particles Protons and electrons are the only particles that have a charge. Protons and neutrons are very close in mass. The mass of an electron is 1/1837 of a proton or neutron. Table 2.1

X H H (D) H (T) U Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei Mass Number X A Z Element Symbol Atomic Number H 1 H (D) 2 H (T) 3 U 235 92 238 2.3

Isotopes of Hydrogen Link to activity

Atomic mass is the mass of an atom in atomic mass units (amu) Micro World atoms & molecules Macro World grams Atomic mass is the mass of an atom in atomic mass units (amu) By definition: 1 atom 12C “weighs” 12 amu On this scale 1H = 1.008 amu 16O = 16.00 amu 3.1

Atomic Mass Atomic and molecular masses can be measured with great precision with a mass spectrometer. 1 amu = 1/12 carbon-12 Link to Movie Figure 2.13

Heavy Light 3.4

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

Average atomic mass of lithium: Natural lithium is: 7.42% 6Li (6.015 amu) 92.58% 7Li (7.016 amu) Average atomic mass of lithium: 7.42 x 6.015 + 92.58 x 7.016 100 = 6.941 amu 3.1

Average atomic mass (6.941)

Do You Understand Isotopes? How many protons, neutrons, and electrons are in C 14 6 ? 6 protons, 8 (14 - 6) neutrons, 6 electrons How many protons, neutrons, and electrons are in C 11 6 ? 6 protons, 5 (11 - 6) neutrons, 6 electrons 2.3