Chapter 4

Atomic Structure

  1. Defining the Atom

An atom is the smallest particle of an element that retains its identity in a chemical reaction.

  1. Early philosophers believed that atoms were indivisible and indestructible.
  2. Dalton’s Atomic theory. Dalton used experimental methods, to transform Democritus’s ideas on atoms into scientific theory.
  3. All elements are composed of tiny particles called atoms.
  4. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.
  5. Atoms of different elements can physically mix together or can chemically combine in simple whole number ratios to form compounds.
  6. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.
  7. Sizing up the Atom

Despite their small size, individual atoms are observable with instruments such as a scanning tunneling microscope.

  1. Structure of the Nuclear Atom

Atoms are divisible and are composed of subatomic particles.

  1. Electrons are negatively charged particles. Electrons are located outside the nucleus of the atom. Electrons have very little mass when compared to the mass of protons and neutrons. Electrons are represented by e-. Electrons have a mass of 9.11 x 10-28 g.
  2. Protons are positively charged particles and are located inside the nucleus. Protons have a mass of 1.67 x 10-24 g. This means that a proton is about 1840 times the mass of an electron. Protons are represented by p+.
  3. Neutrons are neutral particles and are located inside the nucleus. Neutrons have no effective charge and are therefore considered to be neutral. Neutrons have a mass of 1.67 x 10-24 g. Neutrons are represented by n0.
  4. Atomic Nucleus
  5. Most of the volume in an atom is occupied by electrons. Rutherford’s Gold foil experiment proved that the nucleus is very small relative to the total volume of the atom.
  6. The nucleus is the central core of an atom and it is composed of protons and neutrons.
  1. Distinguishing Among Atoms
  2. The number of protons present in an atom determines which atom you have. For this reason the number of protons is also called the atomic number.
  3. The mass of an atom is determined by the number of protons and neutrons. Example: if you have Carbon with six protons and six neutrons the mass number would be 12. The number of neutrons in an atom is the difference between the mass number and the atomic number.
  4. Not all elements are the same. Since the number of protons controls which element you have, what happens when you have a different number of neutrons? Isotopes are what we say you have when you have the same element with a different number of neutrons.

Example:ProtonsNeutrons Mass

Carbon6612

Carbon6713

Carbon6814

In this case all three are Carbon atoms, but they have different masses and are called isotopes.

  1. The mass system for atoms is called an atomic mass unit (amu). This system of mass based on the mass of Carbon-12. An amu is equal to exactly on twelfth the mass of Carbon-12.
  2. A proton has a mass of 1 amu.
  3. A neutron has a mass of 1 amu.
  4. An electron has a mass of 0 amu.
  5. The mass for the elements listed on the periodic table takes account all of the naturally occurring isotopes for an element and their relative abundances. This is called the weighted mass average.

Example:

Carbon has two main isotopes

Abundance

126C 98.89%

136C 1.11%

To solve the weighted take the mass times its abundance

(12).9889+(13)0.0111 = 12.0111

Isotope Problems

  1. ChlorineAbundance

3517Cl 75.77%

3717Cl 24.23%

35.34846

Solve for the weighted mass average.

  1. 63Li 7.5%

73Li 92.5%

6.925

Solve for the weighted mass average.

  1. Boron

105B 19.6%

115B 80.4%

10.804

Solve for the weighted mass average.

  1. Magnesium3 isotopes

2412Mg 78.70%

2512Mg 10.13%

2612Mg 11.17%

24.3247

Solve for the weighted mass average.

  1. 2814Si 92.21%

2914Si 4.70%

3014Si 3.09%

28.10808