Chapter 4

The structure of the atom

Atom

• Smallest part of an element that retains the properties of the element

Democritus:Greek (460-370 BC)

• Proposed the first atomic theory

• called the tiny individual particles “atomos”

• Aristotle said that he was wrong

John Dalton: Eng (1766-1844)

• School teacher• 1808- proposed the

first accepted atomic theory

Joseph John (J.J) Thomson

• English physicist (1856-1940)

• 1897- used the cathode ray tube experiment to discover the electron

• Called plum pudding model

Robert A. Millikan

• American Physicist-(1868-1953)

• 1909- used the oil droplet experiment to determine the charge of an electron and calculate the mass of the electron

• 1923- Nobel Prize• Thomson’s and Millikan’s

works was combined to conclude there must be a positive particle- equal in charge, but more massive than the electron

Ernest Rutherford

• New Zealander (1871-1937)

• 1911- Gold Foil experiment- proved the atom was mostly empty space

• Concluded there was small positive center and called it the nucleus

• “discovered” and named the nucleus

• 1908- Nobel Prize

James Chadwick

• English (1891-1974)• 1932- discovered the

neutron

• Subatomic particles= electron, proton, neutron

• Nuclear forces- short range forces that hold nuclear particles together

• Atomic number = number of protons in an atom– In a neutral atom = # of

electrons

Mass number

• Sum of the protons and neutrons in a nucleus

• Mass number – atomic number = neutrons

Average atomic mass

• Weighted average of the atomic masses of the naturally occurring isotopes of an element

• # on chart• (%)(mass)+ %(mass)= average mass

Example: Carbon

mass number exact weight

percent abundance

12 12.00 98.90

13 13.00 1.10

(12.00) (0.98) + (13.00) (0.01) = 12.011

isotope

• Same element, same number protons, same number electrons, different number of neutrons

Methods of writing isotopes

Nuclear form and hyphen form

Nuclear form=Mass

atomic #

126 C

Hyphen= name-massEx. Carbon-12, carbon-14

Atomic mass unit (AMU)

• 1/12 of the mass of a C-12 atom

• Not exactly 1 proton or 1 neutron

Nuclear reactions

• Reactions which involve as change in an atoms nucleus

Radioactivity

• Substances spontaneously emitting radiation

• Radiation- rays and particles

• Radioactive substances decay until they become stable

4 types of radiation

2.Beta particle- fast moving electron

3. Gamma rays- high energy radiation4.Positron- same mass as electron with (+) charge

Proton decays to a positron and a neutron

1.Alpha particle- a helium nucleus 2 protons and 2 neutrons

0

-1e

4

2He

The Mole

• Equal to the number of particles in exactly 12g of carbon-12

3 equivalents of 1 mole

1. Molar mass: gram atomic mass, formula mass, molecular mass

• Mass number from chart• Add for compounds

Calculate the molar mass of Al(NO3)3

(1 x 27) + (3 x 14) + (9 x 16) = 213.00 g/mol213.00 grams is the mass of one mole of aluminum nitrate.

213.00 grams of aluminum nitrate contains 6.02 x 10^23 entities of Al(NO3)3

2. Avogadro’s number of representative particles

• 6.022 x 1023

• Elements = atoms• Ionic = formula units• Covalent = molecules

Ex: One mole of donuts contains 6.022 x 1023 donuts

3. 22.4 L of a gas at STP

• Standard temp= 0o C, 273 K

• Standard pressure= 1atm, 760 mmHg, 760 torr, 101.325 kPa