Accelerated Chemistry Chapter 3 Notes Mr. Seidel

(Student’s edition)

Chapter 3 problem set:2, 7, 10, 13, 17ace, 18, 20, 21, 22ab, 23ab

3.1The Atom: From Idea to Theory

Historical Background- In approximately 400 BC, Democritus (Greek) coins the term

“atom” (means ). Before that matter was thought to be one continuous piece - called the continuous theory of matter. Democritus creates the discontinuous theory of matter. His theory gets buried for thousands of years

18th century - experimental evidence appears to support the idea of atoms.

Law of Conservation of Mass - Antoine Lavoisier (French) - 1770’s

Law of Definite Proportions - 1799 - Joseph Proust (French) - “The proportions of masses of chemicals in reactions is always the same.”

example

Law of multiple proportions - 1803 - John Dalton - English school teacher “The mass of one element combines with masses of other elements in simple, whole # ratios.”

example

Dalton’s Atomic Theory - Dalton put together the laws of conservation of mass, definite proportion, and multiple proportion to create his own atomic theory.

1. All matter is made up of atoms which are indivisible

*2 Atoms of the same element are identical

*3. Atoms cannot be created or destroyed

4. Atoms combine in simple whole number ratios

5. Chemical reactions are the result of the separating or combining of atoms

* #2 is wrong because of the existence of

* #3 is wrong as an exception is ( )

Page 1

3.2 The Structure of the Atom

Updating Atomic Theory

1870’s - English physicist William Crookes - studied the behavior of gases in vacuum tubes (Crookes tubes - forerunner of picture tubes in TVs). Crookes’ theory was that some kind of radiation or particles were traveling from the cathode across the tube. He named them .

20 years later, J.J. Thomson(English) repeated those experiments and devised new ones. Thomson used a variety of materials, so he figured cathode ray particles must be fundamental to all atoms. 1897 - discovery of the ______.

Charge and Mass of the electron - Thomson and Milliken (oil drop experiment) worked together to discover the charge and mass of the electron

charge = this is the smallest charge ever detected

mass = this weight is pretty insignificant

1909 - Gold Foil Experiment (Rutherford - New Zealand)

Page 2

atom -

atoms are made up of ______,______,______

Subatomic particle / Location / Charge / Relative Mass / Determines…
Proton
Neutron
Electron

Not all atoms of the same element are identical…

Isotope –

. Some isotopes occur naturally - most are produced artificially.

The isotopes of Hydrogen

Name

Drawing

Relative Mass

Relative

Abundance

3.3Weighing and Counting Atoms

We look to the periodic table to give us information about the number of particles are in atoms and also to help us count atoms in a sample.

Page 3

atomic # (Z) - (always a whole number, smaller number on the periodic table) = # of protons in the nucleus - also indicates the # of electrons if the element is not charged

atomic mass – the average mass of all of the isotopes of an element – is a number with a decimal – is always the larger number on the periodic table.

mass number (A) - sum of the protons and neutrons in a nucleus

this number is rounded from atomic mass due to the fact that there are isotopes

# neutrons = example - # of neutrons in Li =

Ion – a charged atom. Atoms become charged by gaining electrons (become a negative charge) or losing electrons (become a positive charge)

Lots of practice:

Element (ion/isotope) / Atomic Number / Atomic Mass / Atomic Mass Number / Proton / Electron / Neutron
He
Be
F
F-1
K
K+1
35Cl
37Cl
25Mg+2

Modern standard of atomic mass and determining masses from weighted averages

Mass of Cl thought to be 35.5 X that of hydrogen

today we know this isn’t true- it’s the weighted average of 2 isotopes

75% 35 Cland 25% 37 Cl

average atomic mass = (%)(mass of 1st isotope)+(%)(mass of 2nd isotope)......

sample problem - find the average atomic mass of B

B11 = 80.20%B10 = 19.80%

Page 4

sample problem - find the %’s of 2 isotopes of Carbon given the following information:

average atomic mass = 12.0111 isotope 1 = 12 C , isotope 2 = 13 C

History lesson - originally H was the basis of all atomic masses and was given the mass of 1.0. Later, chemists changed the standard to oxygen being 16.000 (which left H = 1.008). In 1961, chemists agreed that 12C is the standard upon which all other masses are based.

1/12 of the mass of 1 atom of 12C = 1 amu

The Mole, Avogadro’s number, and Molar Mass

Atoms are tiny, so we count them in “bunches” - a mole is a “bunch of atoms”

mole (definition) - The amount of a compound or element that contains 6.02 x 1023 particles of that substance.

1 mole = 1 molar mass = 6.02 x 1023 particles

Molar Mass - the sum of the atomic masses of all atoms in a formula

Round to the nearest tenth! (measured in amu or grams)

ex - H2H2OCa(OH)2

molar mass is a term that can be used for atoms, molecules (covalent compounds or elements) and formula units (ionic compounds)

Official names may also be:

Formula mass (ionic compounds)

Molecular mass (covalent compounds and diatomic elements)

Atomic weight, Atomic mass, grams formula weight, etc.

examples:1 mole Na = atoms = g

1 mole O2 = molecules = g

1 mole HCl = molecules = g

1 mole NaCl =formula units = g

Page 5

Mole Relationships

For starters in chemistry, we have to be able to convert between moles, grams, and molecules/atoms of substance (also liters when we work with gases)

use the“mole map”

_____ mole

_____ atom/molecule

_____ gram

_____ liter (at STP)

2 steppers

convert 13.8 g Li to moles

convert 2.0 moles Ne to g

convert 3.0 moles of Be to atoms

convert 44.8 L of O2 to moles

3 and 4 steppers

convert 1.2 x 1024 atoms of Magnesium to grams

convert 128 g of O2 to molecules of O2

convert 128 g of O2 to atoms of oxygen

Page 6