The History and Structure of Atoms

Ch. 3: Atoms 3.1 Foundations

History • Democritus • named the most basic particle • atom- means “indivisible • Aristotle • didn’t believe in atoms • thought matter was continuous

History • by 1700s, all chemists agreed: • on the existence of atoms • that atoms combined to make compounds • Still did not agree on whether elements combined in the same ratio when making a compound

Law of Conservation of Mass • mass is neither created or destroyed during regular chemical or physical changes

Law of Definite Proportions • any amount of a compound contains the same element in the same proportions by mass No matter where the copper carbonate is used, it still has the same composition

Law of Multiple Proportions • applies when 2 or more elements combine to make more than one type of compound • the mass ratios of the second element simplify to small whole numbers

Law of Multiple Proportions

Dalton’s Atomic Theory • All mass is made of atoms • Atoms of same element have the same size, mass, and properties • Atoms can’t be subdivided, created or destroyed • Atoms of different element combine in whole number ratios to make compounds • In chemical reactions, atoms are combined, separated, and rearranged.

Modern Atomic Theory • Some parts of Dalton’s theory were wrong: • atoms are divisible into smaller particles (subatomic particles) • atoms of the same element can have different masses (isotopes) • Most important parts of atomic theory: • all matter is made of atoms • atoms of different elements have different properties

Ch. 3: Atoms 3.2 Structure of Atom

Structure of Atom • Nucleus: • contains protons and neutrons • takes up very little space • Electron Cloud: • contains electrons • takes up most of space

Subatomic Particles • includes all particles inside atom • proton • electron • neutron • charge on protons and electrons are equal but opposite • to make an atom neutral, need equal numbers of protons and electrons

Subatomic Particles • number of protons identifies the atom as a certain element • protons and neutrons are about same size • electrons are much smaller • nuclear force- when particles in the nucleus get very close, they have a strong attraction • proton + proton • proton + neutron • neutron + neutron

Atomic Radius • size of atom • measured from center of nucleus to outside of electron cloud • expressed in picometers (1012 pm = 1 m) • usually 40-270 pm

Example • An atom has a radius of 140 pm. How large is that in meters?

Ch. 3 Atoms 3.2b: Important Discoveries

Discovery of Electron • resulted from scientists passing electric current through gases to test conductivity • used cathode-ray tubes • noticed that when current was passed through a glow (or “ray”) was produced

Discovery of Electron Noted Qualities of Ray Produced: • existed- there was a shadow on the glass when an object was placed inside • had mass- the paddle wheel placed inside, moved from one end to the other so something must have been “pushing” it

Discovery of Electron Noted Qualities of Ray Produced: • negatively charged- the rays behaved the same way around a magnetic field as a conducting wire • negatively charged- were repelled by a negatively charged object

Discovery of Electron • all of these led scientists to believe there were negatively charged particles inside the cathode ray • Milliken found the mass of the electron

Discovery of Electron • J.J. Thomson (English 1897) did more experiments to actually make the discovery • he found ratio of charge of this particle to this mass of the particle • since the ratio stayed constant for any metal that contained it, it must be the same in all of the metals

Are electrons the only particles? • since atoms are neutral, something must balance the negative charge • since an atom’s mass is so much larger than the mass of its electrons, there must be other matter inside an atom

Discovery of Nucleus • Rutherford discovered the nucleus by shooting alpha particles (have positive charge) at a very thin piece of gold foil • he predicted that the particles would go right through the foil at some small angle

Discovery of Electron

Discovery of Nucleus • some particles (1/8000) bounced back from the foil • this meant there must be a “powerful force” in the foil to hit particle back Predicted Results Actual Results

Discovery of Nucleus Characteristics of “Powerful Force”: • dense- since it was strong enough to deflect particle • small- only 1/8000 hit the force dead on and bounced back • positively charged- since there was a repulsion between force and alpha particles

Find the element Sodium on your periodic table. What do you know about atoms of sodium from the information on the table?

Ch. 3 Atoms 3.3 Counting Atoms

Atomic Number • (Z) number of protons • All atoms of the same element have the same atomic number • located above the symbol in the periodic table • order of the elements in the periodic table

Isotopes • atoms of the same element with different numbers of neutrons • most elements exist as a mixture of isotopes • What do the Carbon isotopes below have in common? What is different about them?

Mass Number • sum of particles in nucleus • A = #p + #n Hydrogen isotopes have special names: • protium • deuterium • tritium • What do the prefixes in their names come from?

Designating Isotopes • Hyphen notation: • Name - mass number • ex. Carbon – 13 • Nuclear Symbol notation:

Examples • 7 protons, 8 neutrons Nitrogen-15 • 17 electrons, 19 neutrons Chlorine- 36

Examples • Z=5, 6 neutrons Boron- 11 • A=75, 42 neutrons Arsenic- 75

Examples 6 7 6 6 13 54 77 54 54 131 11 13 11 11 24 8 7 8 8 15

A neutral atom contains 34 electrons and has an A of 59. Write the nuclear symbol notation and hyphen notation for this isotope.

Ch. 3 Atoms 3.3 Counting Atoms

Relative Atomic Mass since masses of atoms are so small, it is more convenient to use relative atomic masses instead of real masses to set up a scale, we have to pick one atom to be the standard since 1961, the carbon-12 nuclide is the standard and is assigned a mass of exactly 12 amu

Relative Atomic Mass atomic mass unit (amu)- one is exactly 1/12th of the mass of a carbon-12 atom mass of proton= 1.007276 amu mass of neutron= 1.008665 amu mass of electron= 0.0005486 amu

Relative Atomic Mass the mass number (A) and the relative atomic mass are very close but not the same because relative atomic mass includes electrons the proton and neutron masses aren’t exactly 1 amu

Average Atomic Mass weighted relative atomic masses of the isotopes of each element each isotope has a known natural occurrence (percentage of that elements’ atoms)

Calculating Average Atomic Mass Naturally occurring copper consists of: 69.71% copper-63 (62.929598 amu) 30.83% copper-65 (64.927793 amu) (0.6971 x 62.929598)+(0.3083 x 64.927793) =63.55 amu

Calculating Average Atomic Mass An element has three main isotopes with the following percent occurances: #1: 19.99244 amu, 90.51% #2: 20.99395 amu, 0.27% #3: 21.99138 amu, 9.22% Find the average atomic mass and determine the element.

Calculating Average Atomic Mass

The mole a unit for measuring a very large amount- like number of atoms or molecules in a sample like one dozen (1 dozen = 12 things) except bigger: 1 mole = 6.022x1023 things Why 6.022x1023 ? 6.022x1023 is the number of atoms in exactly 12 g of carbon-12

The mole 6.022x1023 is called Avogadro’s Number in honor of all of his contributions to chemistry can be used as a conversion factor between a number of things and mole

Molar Mass the mass of one mole of pure substance in grams per mole numerically equal to average atomic mass under the symbol on the periodic table can be used as a conversion factor between moles and grams

Conversion Factors # Atoms Grams Moles Use Molar Mass: grams per mole Use Avog.’s Number: atoms per mole

The History and Structure of Atoms