chapter 3 – atoms: the building blocks of matter 3.1: atomic theory history a. 1700s: quantitative...

Chapter 3 – Atoms: The Building Blocks of Matter

• 3.1: Atomic Theory History A. 1700s: quantitative studies of chemical reactions led to several laws: 1. law of conservation of mass – mass is

neither created nor destroyed in a chemical reaction

2. law of definite proportions – a chemical compound always contains the same elements in the same proportion by mass

ex. NaCl is always 39.3% sodium and 60.7% chlorine by mass.3. law of multiple proportions – if the same 2 elements are found in different compounds, then the ratio of the masses of the second element (with the first element’s mass being the same) is always a whole number

B. 1800s: Dalton’s Atomic Theory –He explained the above 3 laws in his theory:

• 1. All matter is composed of atoms• 2. Atoms of a given element are

identical. Atoms of different elements are different.

• 3. Atoms cannot be subdivided, created, or destroyed

• 4. Atoms of different elements combine in whole-number ratios

• 5. Atoms are combined, separated, or rearranged in a chemical reaction

C. The Modern Atomic Theory – A couple of Dalton’s points have been modified:• Atoms are divisible• Atoms of the same element can have different

masses.

• 3.2: The Structure of the AtomA. The Electron – discovered in 1897 by J.J. Thomson after experiments with a cathode-ray tube. Properties:

negatively charged 1/1837 the mass of a proton symbol = e-

B. The Nucleus – discovered by Rutherford after doing his Gold Foil Experiment. Composition of the nucleus:

1. protons = positively charged subatomic particles2. neutrons = subatomic particles with no charge

• Both protons and neutrons have a mass of 1amu (atomic mass unit)

• Rutherford's experiment

3.3: Counting AtomsA. Atomic Number = the # of protons• Each element has its own atomic #• In a neutral atom, the # of protons = the # of e-

B. Mass Number = the # of protons + the # of neutrons in an atom (e- do not contribute significantly to an atom’s

mass because their mass is too small)C. Isotopes = atoms of the same element (same # of protons) that have different numbers of neutrons.• 2 ways of writing:• 1. element name – mass # (ex. Hydrogen - 3)• 2.

D. Average Atomic MassMost elements exist in nature as a mixture of isotopes. Average atomic mass is the weighted average mass of all the isotopes of an element. In calculating atomic mass, we must consider the abundance of each isotope.– Steps in calculating atomic mass:

1. Multiply the mass of each isotope by the relative abundance of that isotope

2. Total the answers from step #1

Chapter 21 – Nuclear Chemistry

21.2: Radioactive DecayA. Terms:

• radioactive decay = the process by which an unstable nucleus loses energy by emitting penetrating rays called radiation• radioisotope (or

radioactive nuclide) = an isotope of an element that has an unstable nucleus

B. Types of Radiation:1. alpha (α): 2 protons + 2 neutrons (same thing as a helium

nucleus)• -not very penetrating (can be stopped by paper or skin)• Results in the atomic number going down by 2 and the mass

going down by 4

2. beta (β): electrons formed from the decomposition of a neutron. The proton formed stays in the nucleus (atomic number inc. by 1)

-more penetrating than alpha (can be stopped by wood or foil)

3. gamma (γ): high-energy electromagnetic radiation. Often emitted along with alpha or beta.

• -has no mass or charge• -very penetrating (need concrete or

thick lead to stop

21.1: Nuclear Stability• The stability of a nucleus

depends on its neutron to proton ratio.

• For atoms with an atomic # < 20, a 1:1 ratio is stable.

• For atoms with an atomic # > 20, more neutrons are needed for stability, and a ratio of 1.5:1 is most stable for heavy elements

• Note: all nuclei with atomic # >83 are radioactive (not stable)

• Too many neutrons leads to beta decay, which results in a decrease in neutrons and an increase in protons (atomic # increases by 1)

• A nucleus with too many protons and neutrons will emit alpha particles (atomic # decreases by 2)

21.2: Half-Life = the time required for half of a radioactive sample to decay

– ex. half life of uranium-238 is 4.5 x 109 years– Half-life problems: see sample problem B on page 689

•

• 3.3: Counting AtomsA. The Mole = 6.022 x 1023 particles of something (also

called Avogadro’s number)B. Molar mass = the mass of one mole of a substance. The

units are g/mole. It’s numerically equal to the atomic mass in atomic mass units. Sometimes it’s called the formula mass.

• For ex., the molar mass of sodium is 23.0g/mole. Thus, 23g of sodium = 1 mole = 6.022 x 1023 atoms.

C. Conversions:• between mass and moles of an element: use the molar

mass as a conversion factor• between particles and moles of an element: use

6.022 x 1023 particles/mole as a conversion factor

• See sample problems B-E on pages 84-86