Chapter 3 Notes

The Atom: From Idea to Theory

Historical BackgroundIn approximately 400 BC, Democritus (Greek)

coins the term“atom” (means indivisible). 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 Lavosier (French) -1700’s The number of each kind of atoms on the

reactant side must equal the number of each kind of atoms on the product side

A + B + C ABC

Law of Multiple Proportions

John Dalton (English) - 1803 The mass of one element combines with

masses of other elements simple in whole number ratios.

Water (H2O) is always: 11.2% H; 88.8% O

Sugar (C6H1206) is always: 42.1% C; 6.5% H; 51.4% O

Law of Multiple Proportions Cont’d

Wt. of H

Wt. of O

H + O H2O 2 16

H + O H2O2 2 32

The ratio of O in H2O2 to O in H2O = _______2 to 1______________

Dalton’s Atomic Theory

1. Everything is made of atoms.

2. Atoms of the same element are identical. (NOT)

3. Atoms can not be broken down, created, or destroyed. (NOT)

4. Atoms combine in simple whole number ratios to form chemical compounds. 2:1 1:1

5. A chemical reaction is the combining, separation, or rearrangement of atoms.

C (s) + O2 (g) --------> CO2 (g)

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 electrons

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 positive charge

Plum Pudding Model

The Structure of the Atom Cont’d

o Charge and Mass of the electron o Thomson and Milliken (oil drop experiment) worked

together to discover the charge and mass of the electron

charge = 1.592 × 10−19 coulombs this is the smallest charge ever detected

mass = 9.1093821545 × 10−31 g this weight is pretty insignificant

The Structure of the Atom Cont’d

1909 - Gold Foil Experiment (Rutherford - New Zealand)

Nuclei are composed of ‘nucleons’: protons and neutrons

Top: Expected results: alpha (+) particles passing through the plum pudding model of the atom undisturbed.Bottom: Observed results: a small portion of the particles were deflected, indicating a small, concentrated positive charge.

Important Subatomic Particles

a.m.u. Mass, kg Charge Location

Job

Proton (p+)

1 1.67265×10-27 +1 nucleus

ID

Neutron (n°)

1 1.67495×10-27 0 nucleus Stabilize atom

Electron (e-)

0 9.10953×10-31 -1 clouds Bonding

Important Subatomic Particles Cont’d

Electrostatic force - pulls nuclei apart: protons and neutrons

Strong Nuclear Force- force holds nuclei together

Weighing 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.

Counting Atoms Atomic Number (Z)

Number of protons in the nucleusUniquely labels each element

Mass Number (M) Number of protons + neutrons in the nucleus

Weighing and Counting Atoms

Counting electrons Atoms

Same number of electrons and protons Ions – lost or gained electrons Ionic charge (q) = #protons - #electrons

Positive ions are cationsNegative ions are anions

Weighing and Counting Atoms

If the mass # comes from the p+ and n0 [each with masses of exactly 1], why don’t the atomic weights/masses of the all elements turn out to be whole numbers?

Because the atomic weights/masses on the P-Table are the “weighted averages,” of the naturally occurring isotopes of the element. (remember: ignore the mass of the e-, it’s too small to care about.

Review of Formulas

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 = A - Z example - # of neutrons in Li = 6.941-3 = 3.941 rounds to 4

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

Lets Practice!

p+ e- n° Atomic # =(# of p+)

Mass # =(p+ + n0)

C 6 6 6 6 12

Ca 20 20 20 20 40

U 92 92 146 92 238

Cl 17 17 18 17 35

Mg 12 12 12 12 24

14C 6 6 8 6 14

S-2 16 18 16 16 32

Na+1 11 10 12 11 23

Isotopes

Two atoms of the same element (same # of p+) but with different weights (different # of n0)

Average Atomic Mass (“weighted average”) Definition - The average weight of the natural isotopes of an

element in their natural abundance.

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 12 - C is the standard upon which all other masses are based.

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

Isotope Calculations

Carbon consists of two isotopes: 98.90% is C-12 (12.0000 amu). The rest is C-12 (13.0034 amu). Calculate the average atomic mass of carbon to 5 significant figures.

12.011 amu

Chlorine consists of two natural isotopes, 35Cl (34.96885) at 75.53% abundance and 37Cl (36.96590) at 24.47% abundance. Calculate the average atomic mass of Chlorine.

35.46

Antimony consists of two natural isotopes 57.25% is 121Sb (120.9038). Calculate the % and mass of the other isotope if the average atomic mass is 121.8.

The Mole, Avogadro’s number and Molar Mass

The MoleAtoms are tiny, so we count them in “bunches”.A mole is a “bunch of atoms”.The Mole (definition) -The amount of a

compound or element that contains 6.02 x 1023 particles of that substance.

1 mole = 1 gram formula mass = 6.02 x 1023 particles

The Mole, Avogadro’s number and Molar Mass

Molar Mass 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 - H2 H2O Ca(OH)2

2.0g 18.0 g 74.1 g

The Mole, Avogadro’s number and Molar Mass

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.

The Mole, Avogadro’s number and Molar Mass

Examples: 1 mole Na = 6.02 x 1023 atoms = 23.0 g

1 mole O2 = 6.02 x 1023 molecules =

1 mole HCl = 6.02 x 1023 molecules =

1 mole NaCl =6.02 x 1023 formula units=

Mole Map

Liters

Atoms, molecules, particles

Grams

Mole

22.4 L

6.022 x10 23Molar Mass

Examples

2 steppersconvert 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

Examples

3 and 4 steppersconvert 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

Convert 100. g of Ar to liters of Ar