Unit 3 – Atomic Structure

Bravo – 15,000 kilotons

Democritus

• 400 BC

• Greek philosopher

• 1st to come up with idea of atoms

John Dalton – 1800’s

• Major contributor of Atomic Theory1)All matter made of atoms2)All atoms of an element are alike3)Atoms cannot be created or

destroyed4)Atoms combine in whole-number

ratios to form compounds

JJ Thomson – late 1800’s

• Cathode Ray Experiment – discovery of electrons

• “Plum Pudding” model of atom

• Measured charge to mass ratio of e-

Thomson’s Atomic Model

Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

Discovery of the ElectronDiscovery of the ElectronI n 1897, J .J . Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.

Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

Conclusions from the Study of the Electron

Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.

Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons

Electrons have so little mass that atoms must contain other particles that account for most of the mass

Millikan - 1910

Oil Drop Experiment

Determined actual charge and mass of an e-

Rutherford - 1910

Discovered nucleus and that it was positive

Gold Foil Experiment1) Most of atom is empty space (majority of particles went straight through)

2) nucleus is small, dense and positively charged (some positive charges were greatly deflected)

Rutherford’s Gold Foil Experiment

Alpha particles are helium nuclei Particles were fired at a thin sheet of

gold foil Particle hits on the detecting screen

(film) are recorded

Rutherford’s Findings

The nucleus is small The nucleus is dense The nucleus is positively charged

Most of the particles passed right through

A few particles were deflected VERY FEW were greatly deflected

“Like howitzer shells bouncing off of tissue paper!”

Conclusions:

Niels Bohr - 1915

• Proposed early model of atom• “Planetary Model” electrons orbit

nucleus like planets orbit sun• Lacks math of modern version• Has some errors/violates current theory • Radiation is emitted when electrons

move from one orbit to another

Chadwick

Discovered neutrons

Modern Atomic Theory(changes from Dalton)

Atoms of an element have a characteristic average mass which is unique to that element (isotopes)

Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

Foundations of Atomic Theory

Law of Conservation of Mass: mass is neither created or destroyed in ordinary chemical reactions

Law of Definite Proportions (composition): compounds contain same elements in same ratio by massExample: NaCl is always 39.9% Na and 60.66% Cl by mass

Law of Multiple Proportions: 2 or more different compounds composed of same two elements have ratios of small whole numbersExample: CO vs CO2 ratio of oxygen to oxygen is 2 to 1

What is AMU?

Stands for atomic mass unit – used when describing “relative” atomic masses

This system is used because the actual masses of atoms are so small

Carbon-12 is the standard to which all other elements are compared (i.e. hydrogen-1 has a mass that is 1/12 that of carbon-12 so it’s mass would be 1 amu)

Atomic NumberAtomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.

Element # of protons Atomic # (Z)

Carbon 6 6

Phosphorus 15 15

Gold 79 79

Mass NumberMass number is the number of protons and neutrons in the nucleus of an isotope.Mass # = p+ + n0

Nuclide p+ n0 e- Mass #

Oxygen - 10

- 33 42

- 31 15

8 8 1818

Arsenic 75 33 75

Phosphorus 15 3116

Atomic Masses

Isotope Symbol Composition of the nucleus

% in nature

Carbon-12

12C 6 protons6 neutrons

98.89%

Carbon-13

13C 6 protons7 neutrons

1.11%

Carbon-14

14C 6 protons8 neutrons

<0.01%

Atomic mass is the average of all the naturally isotopes of that element.Carbon = 12.011

IsotopesIsotopes are atoms of the same element having different masses due to varying numbers of neutrons.Isotope Proto

nsElectron

sNeutron

sNucleus

Hydrogen–1

(protium)

1 1 0

Hydrogen-2

(deuterium)

1 1 1

Hydrogen-3

(tritium)

1 1 2

ISOTOPES

• Example: • Carbon, C, exists in 3 isotopes:

Isotope Protons Neutrons

Mass

Carbon-12

6 6 12

Carbon-13

6 7 13

Carbon-14

6 8 14

Isotope symbols

Hyphen notation Nuclear notation

Carbon – 12 12C 6p+ and 6no

Carbon – 13 13C 6p+ and 7no

Carbon – 14 14C 6p+ and 8no

6

6

6

How to Calculate the Average Mass

What is the average atomic mass of sample of Cesium with 3 isotopes:

75% 133Cs, 20% 132Cs, and 5% 134Cs. 0.75 x 133 = 99.750.20 x 132 = 26.400.05 x 134 = 6.70

Total = 132.85 avg. atomic mass

The Mole

1 dozen =1 gross =

1 ream =

1 mole =

12

144

500

6.02 x 1023

There are exactly 12 grams of carbon-12 in one mole of carbon-12.

Avogadro’s Number6.02 x 1023 is called “Avogadro’s Number” in honor of the Italian chemist Amadeo Avogadro (1776-1855).

Amadeo Avogadro

I didn’t discover it. Its just named after

me!

Calculations with Moles:Converting moles to grams

How many grams of lithium are in 3.50 moles of lithium?

3.50 mol Li= g Li

1 mol Li

6.94 g Li45.1

Calculations with Moles:Converting grams to moles

How many moles of lithium are in 18.2 grams of lithium?

18.2 g Li= mol Li

6.94 g Li

1 mol Li2.62

Calculations with Moles:Using Avogadro’s Number

How many atoms of lithium are in 3.50 moles of lithium?

3.50 mol Li = atoms Li

1 mol Li

6.022 x 1023 atoms Li 2.11 x 1024

Calculations with Moles:Using Avogadro’s Number

How many atoms of lithium are in 18.2 g of lithium?

18.2 g Li

= atoms Li

1 mol Li 6.022 x 1023 atoms Li

1.58 x 1024

6.94 g Li 1 mol Li

(18.2)(6.022 x 1023)/6.94