Atomic Theory Chapter 3 Sections 1 &2 9/18/14.

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Presentation transcript:

Atomic Theory Chapter 3 Sections 1 &2 9/18/14

WHAT’S A THEORY?

Atomic Theory

The Ancient Greeks Democritus and other Ancient Greeks were the first to describe the atom around 400 B.C. The atom was nature’s basic particle that makes up all matter. The word atom comes from the Greek word “atomos” meaning indivisible.

Aristotle The Greek philosopher Aristotle did not believe in atoms, and instead believed all matter was continuous. Aristotle’s belief was accepted for ~2000 years, as there was no evidence to support either theory.

Antoine Lavoisier First to state the law of conservation of mass, which states that mass is neither created nor destroyed during chemical or physical changes.

Joseph Louis Proust Law of Definite Proportions = chemical compounds contain the same elements in the same proportions by mass regardless of the source or the amount. i.e. Sodium Chloride (table salt) is always 39.34% Na and 60.66% Cl by mass.

“The BIG 3” 3 Major Theories on the ATOM Dalton Thompson Rutherford

John Dalton’s Atomic Theory All matter is composed of extremely small particles called atoms. Atoms of an element are identical to each other and different from atoms of other elements.

3) Atoms cannot be created or destroyed. 4) Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5) Atoms are combined, separated or rearranged when chemical reactions occur.

Impact of Dalton’s Atomic Theory By relating atoms to mass, Dalton turns Democritus’ ideas into a scientific theory which could be tested. Dalton’s model of the atom is a small, solid, and indivisible particle. (known as the “ball-bearing” model)

Flaws in Dalton’s Theory Not all parts of Dalton’s Theory have proven to be correct. Today we know that… Atoms are divisible into even smaller particles Atoms of the same element can have different masses.

Cathode-Ray Tube Experiments Electric current was passed through gases at low pressures. It was observed that the end of the tube near the cathode glowed and a ray traveled from there to the anode when current was passed through the tube.

In these experiments, it was also observed that the rays were deflected away from magnetic fields and negatively charged objects .

J.J. Thomson Carries out experiments that support the hypothesis that cathode rays are made of negatively charged particles. Measured the charge to mass ratio of cathode ray particles and determined that it was the same regardless of the metal or gas used. Thomson concludes that all cathode rays are composed of identical negatively charged sub-atomic particles, which he calls electrons.

Thomson’s Model of the Atom Thomson’s experiments show that (1) atoms are divisible and (2) they include negatively charged electrons. Thomson’s Plum-Pudding Model of the Atom has negative particles spread evenly throughout a solid, positively charged sphere.

Robert Millikan Conducted an Oil Drop Experiment that measured the charge of the electron. Using this charge and Thomson’s charge to mass ratio, the mass of the electron is also discovered. Mass of 1 electron = 9.109 x10-31 kg

Ernest Rutherford Conducted the Gold-Foil Experiment Positively charged alpha particles were fired at a thin piece of gold foil. It was expected that the alpha particles would pass through with only a slight deflection.

However, it was discovered that about 1 out of every 8,000 particles were actually deflected back toward the source. Rutherford concludes that there must be a very small and densely packed region of positive charge in the atom that would repel these alpha particles. Rutherford is credited with discovering the nucleus.

The Nuclear Model of the Atom Nucleus = very small and dense center of the atom, where protons and neutrons are found. Protons have a positive charge. Neutrons are electrically neutral. Negatively charged electrons orbit the nucleus like the planets orbit the sun. Most of the atom is empty space.

Mass of protons and neutrons = 1 *Mass of protons and neutrons = 1.67 x10-27 kg compared to the electron’s mass of 9.109 x10-31 kg Subatomic Particles Proton Neutron Electron Charge + — Mass* 1 amu ~ 0 Location Nucleus Electron Cloud 21

Atomic Mass Units AMU = Atomic Mass Unit Masses of atoms are very small Carbon-12 atom is used as the standard atom for a mass reference. 1 amu = 1/12 of C-12 atom Protons = 1.007276 amu ~= 1.0 amu Neutrons = 1.008665 amu ~= 1.0 amu Electrons = .0005486 amu ~= 0.0 amu

Average Atomic Mass Most elements occur in nature as a mixture of isotopes

Average Atomic Mass Weighted average of the atomic masses of the naturally occurring isotopes Think about your Warm Up Questions

Average Atomic Mass / Isotopes

Change in Charge (“Ions”) Protons +1 Charge for each Proton Ex: Carbon12 6 protons (+6) 6 Neutrons (no charge) Mass Number = # protons + # Neutrons Electrons -1 Charge for each Electron Ex: Carbon-12 6 electrons (-6) C-12 is a neutral Atom (+6) + (-6) What if add an electron (-1)?