Chapter 3

Atoms: The Building Blocks of Matter

Sections 3-1

Objectives

• Explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions.
• Summarize the five essential points of Dalton’s atomic theory.
• Explain the relationship between Dalton’s atomic theory and the law of conservation of mass, the law of definite proportions, and the law of multiple proportions.

9. The atom: From philosophical Idea to scientific theory.
1. Foundations of Atomic Theory
1. 1700’s elements were known as a substance that cannot be further broken down by ordinary means and can combine to make compounds
2. 1790’s changed due to quantitative data
1. Law of conservation of mass- mass is neither destroyed nor created during ordinary chemical reactions.
2. Law of definite proportions- chemical compounds contain the same proportions by mass regardless of the size of the sample or source.
3. Law of multiple proportions- if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.
1. Which means the little subscript is always whole numbers.
2. Dalton’s Atomic Theory
1. 1808 the atom theory was improved
2. Dalton’s theory has 5 parts
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds
5. In chemical reactions, atoms are combined, separated, or rearranged.
3. Modern Atomic Theory
1. Democritus’s first idea of an atom
2. Dalton’s atomic theory
3. Parts of Dalton’s theory that are wrong
1. Atoms are divisible= nuclear bomb
2. Atoms of the same element can have different number of subatomic particles.

Section 3-2

The structure of the Atom

Objectives

• Summarize the observed properties of cathode rays that led to the discovery of the electron.
• Summarize the experiment carried out by Rutherford and His co-workers that led to the discovery of the nucleus.
• List the properties of protons, neutrons, and electrons
• Define atom.

2. The structure of the atom. Atom the smallest form of matter that still has the same properties

1. Discovery of the electron

1. Late 1800’s thru the use of a cathode ray tube
2. J.J. Thomson given credit
3. Investigators noticed that when current was passed through a cathode-ray tube the opposite end glowed.
4. Led to the hypothesis that the particles where negatively charged.
5. Later named Electrons
6. Charge and Mass of the electron

1. Robert. Millikan
2. Oil drop experiment
3. 9.109 X 10-³¹kg or 1/1837 of a hydrogen atom

7. So. If an atom has extremely low weight particles that are negative? what else is in the atom.

2. Discovery of the Atomic Nucleus

1. 1911
2. Ernest Rutherford, Hans Geiger and Ernest Marsden.
3. Gold foil experiment using alpha particles
4. Proved

1. Most of the atom is empty space
2. A very dense packed area with a positive charge called the nucleus.
3. No proof –that the electrons where surrounded by nucleus.

3. Composition of the atomic Nucleus

1. Nucleus contains two parts

1. Protons

1. Positive charge
2. Mass- 1.673 x 10^-27 kg

2. Neutrons

1. No Charge
2. Mass- same as electron

2. What holds the Nucleus together?

1. Like charges repel (in the nucleus 2 or more protons should repel)
2. Nuclear forces- short-range forces that hold the particles together.

4. The Sizes of Atoms

1. Electron Cloud – area around the nucleus containing the electrons.
2. Use picometer to measure atoms

Section 3-3

Counting Atoms

Objectives

• Explain what isotopes are.
• Define atomic number and mass number, and describe how they apply to isotopes.
• Given the identity of a nuclide, determine its number of protons, neutrons, and electrons.
• Define mole in terms of Avogadro’s number, and define molar mass.
• Solve problems involving mass in grams, amount in moles, and number of atoms of an element.

2. Counting Atoms
1. Atomic Number-
1. Number of protons in an atom
2. The small number in the box
2. Isotopes-
1. Deals with the number of neutrons
2. Atoms can have different numbers of neutrons
3. Are atoms with different masses
4. Ex. Protium, Deuterium, and Tritium are isotopes of hydrogen
5. To calculate the number of neutrons subtract big minus little.
3. Mass Number
1. The total number of protons and neutrons
2. The big number in the box
4. Designating Isotopes
1. Isotopes are called Nuclide
2. Written in two forms
1. Hydrogen-3
5. Practice pg 77
6. Relative atomic masses
1. Atomic mass unit ( amu )
2. 1/12 the mass of a carbon-12 atom
7. Average atomic masses of elements
1. The weighted average of all the isotopes
2. The reason for decimals
3. Calculating Average atomic masses (like grades)
8. Relating mass to numbers of atoms
1. The mole
1. SI unit for the amount of substance present.
2. is the amount of substance that contains as many particles as there are atoms in exactly 12 g of carbon-12.
2. Avogadro’s number
1. Number of particles in one mole of any substance
2. 6.02 X 10²³
3. Molar Mass
1. The mass of one mole of a pure substance
2. Equal to the rounded mass number form the box.
4. Gram to mole conversions
5. Conversions with Avogadro’s number.