Atomic StructureChapter 5

The Atomists: The first atomic theory

• 460 BCE: Greek Democritus suggested that matter is “ composed of minute, invisible, indestructible particles of pure matter which move about eternally in infinite empty” http://www.winneconne.k12.wi.us/middle_school/7th%20Grade/LENZ/history_of_atomic_theory.htm

Development of Modern Atomic Theory

• In 1782, a French chemist, Antoine Lavoisier (1743-1794), made measurements of chemical change in a sealed container.

• He observed that the mass of reactants in the container before a chemical reaction was equal to the mass of the products after the reaction.

Law of Conservation of Mass• Lavoisier concluded

that when a chemical reaction occurs, mass is neither created nor destroyed but only changed.

• Lavoisier’s conclusion became known as the law of conservation of mass.

Dalton’s Atomic Theory• John Dalton (1766-

1844), an English schoolteacher and chemist, studied the results of experiments by Lavoisier and many other scientists.

• Dalton proposed his atomic theory of matter in 1803.

Dalton’s Atomic Theory (The main points)

1. All matter is made up of indivisible particles called atoms.

2. All atoms of one element are exactly alike, but are different from atoms of other elements.3. Atoms can mix in simple whole number

ratios to form compounds.4. Chemical reactions happen when atoms

are separated, joined, or rearranged.

The Electron• Because of Dalton’s atomic theory, most scientists in

the 1800s believed that the atom was like a tiny solid ball that could not be broken up into parts.

• In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model was not accurate.

Cathode-Ray Tube

• Thomson knew that objects with like charges repel each other, and objects with unlike charges attract each other

The Electron

• Thomson named the negative particles found in his experiment electrons.

• Electrons (e-):

– Are the smallest subatomic particles

– Have a charge of -1

Protons

• Atoms aren’t negative, so there must be something positive in the atom to balance the electrons.

• Protons (p+):– Positively charged subatomic particles (+1)– Mass 1 proton = mass 1840 electrons

• At this point, it seemed that atoms were made up of equal numbers of electrons and protons

• However, in 1910, Thomson discovered that neon consisted of atoms of two different masses

Neutrons

• Neutrons (n0):

• Neutral charge

• Same mass as a proton

• Because of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained these differences in mass.

Rutherford’s Gold Foil Experiment

• In 1909, a team of scientists led by Ernest Rutherford in England carried out the first of several important experiments that revealed an arrangement far different from the cookie-dough model of the atom.

The Gold Foil Experiment

The Nuclear Model of the Atom

• Because most of the particles passed through the foil, they concluded that the atom is nearly all empty space.

• To explain the results of the experiment, Rutherford’s team proposed a new model of the atom.

The Nuclear Model of the Atom

• Because so few particles were deflected, they proposed that the atom has a small, dense, positively charged central core, called a nucleus.

• Nucleus:

• Contains most of the mass of the atom (protons and neutrons)

• Is very small compared to the rest of the atom

You know you want to know how many of these particles are in an

atom, right?

• Atomic Number : # of protons in the atom

• Mass number: # of protons + # neutrons

How to get the info from the periodic table

H1

1.0013

Atomic Number: Number of protons in the nucleus of the atom.

In this chapter you will fill out lots of tables…

Name Protons Electrons Neutrons

Helium-4

Carbon-12

Carbon-13

How many subatomic particles?

• # protons = atomic number (always)

• For a NEUTRAL atom # electrons=# protons (this will be a bit different in upcoming chapters…)

• # neutrons = mass number - # protons (always)

In this chapter you will fill out lots of tables…

Name Protons Electrons Neutrons

Helium-4

Carbon-12

Carbon-13

Carbon-12 and Carbon-13 are isotopes

Isotopes

• Atoms of an element that are chemically alike but differ in mass are called isotopes of the element.

So why the difference in mass?

So what is the bottom number for?

H1

1.0013

Atomic Number: Number of protons in the nucleus of the atom.

The atomic mass unit (amu)• Atoms don’t have much mass! (10-23 g)

– This is small, and a bit hard to work with

• Use carbon-12 isotope as a ref. and make 12 carbon atom= 1 amu

• On the periodic table the value is a weighted average of all isotopes of the element – that is why it isn’t a whole number (includes carbon-12, carbon-13, etc).

The Periodic Table • Chemists kept discovering new

elements and decided they needed a way to organize them.

• Dmitri Mendeleev constructed the 1st periodic table– Grouped elements according to

similar chemical properties– He was so smart he left spaces

for elements that would be discovered later on.

The modern periodic table

• Henry Moseley tweaked the table and ordered by atomic number rather than atomic mass.

PeriodGroup or family

Alkali Metals

Alkaline Earth Metals

Transition Metals

Inner Transition Metals

Nonmetals

MetalsHalogens

Noble Gases