John Dalton l English scientist l 1766 1844

John Dalton l. English scientist l 1766 -1844 l. Dalton’s atomic theory

Dalton’s Theory: l 1. All elements are composed of tiny indivisible particles l 2. Atoms of the same element are identical. – The atoms of any one element are different from those of any other element.

Dalton’s Theory cont. l 3. Different elements can physically or chemically combine to form compounds l 4. Chemical reactions occur when elements are separated, joined or rearranged. – We cannot change an atom of one element into an atom of another element by chemical means

The atom: l. The atom composed of the following particles: – Electrons – Protons – Neutrons

Electrons: l. Negatively charged l. J. J. Thomson discovered electrons with cathode ray tube l. Found in “electron cloud” surrounding nucleus of atom

Protons: l. Positively charged sub atomic particles

Neutrons: l. Neutrally charged subatomic particles l. Have a mass nearly equal to protons

Distinguishing between the atoms l. Atomic number = the number of protons l. Mass number = the total number of protons + neutrons l. Number of neutrons = the mass number – number of protons

Table 5. 2 l. Table 5. 2 is a list of the first ten elements and some relative information l. Properties such as atomic number and mass number are identifying features of an element

Examples: l. Practice problems pg. 115 and 116 l. Before we answer these questions lets look at the periodic table of elements.

PCA: l. Pg 114 l. Chem. Math l. Practice problems A-N

Isotopes: l. What is an Isotope? – Hint: think back to the lab with the fictitious element vegium! l. Isotope = same number of protons different number of neutrons – Give an example: l. Carbon 14 (Carbon 12)

Hydrogen Isotopes: l. Hydrogen has three isotopes l. Most common- Hydrogen – 1 proton – No neutrons l. Second isotope = Hydrogen – 2 – 1 proton and 1 neutron – Also called deuterium

Hydrogen Isotopes: l. Third isotope = Hydrogen-3 – 1 proton – 2 neutrons – Also called tritium

Practice Problems: l. Try 117 the practice problems on pg.

Atomic Mass: l. Atomic mass vs. actual mass l. Actual mass is measured with a mass spectrometer l. Atomic mass is determined by a reference to a standard isotope l. Carbon 12 was chosen to be the reference isotope

Atomic Mass: l. Carbon was assigned 12 amu’s – Or atomic mass units – amu’s are defined as one – twelfth the mass of carbon – It was assigned this number because Carbon has 6 protons and 6 neutrons, of which almost all of the mass of the element is composed

How do we explain atomic mass: l. Definition: atomic mass of an element is the weighted average mass of the atoms in a naturally occurring sample of an element. – This reflects both mass and relative abundance

Lets try again: l. Naturally, elements occur as a mixture of two or more isotopes. l. Each isotope having a fixed mass and percent abundance l. The amu of an element is usually very close to the atomic mass number of the most frequently occurring isotope

Calculating atomic mass: l. Need three numbers – 1. The number of stable isotopes of the element – 2. The mass of each isotope – 3. The natural percent abundance of each isotope

Sample Problems: pg 121 l. Review the sample problems then attempt the practice problems on page 121.