Chapter 5 Atomic Structure & the Periodic Table. Early Scientists n As scientists of the...

Chapter 5

Atomic Structure

&

the Periodic Table

Early Scientists As scientists of the eighteenth century studied the

nature of materials, several things became clear:– 1. Most natural materials are mixtures of pure

substances.– 2. Pure substances are either elements or

combinations of elements called compounds.– 3. A given compound always contains the same

proportions of the elements - Law of Constant Composition.

Early Models of the Atom

John Dalton– English teacher– 1766-1844– Studied ratios of elements in chemical

reactions.– Formulated hypotheses and theories to explain

his observations and came up with Dalton’s atomic theory.

Dalton’s Atomic Theory

1. All elements are composed of tiny particles called _________________.

Dalton’s Atomic Theory

2. Atoms of the same element are _________________.

Dalton’s Atomic Theory

3. The _________________ of a given element are different from those of any other element.

Dalton’s Atomic Theory

4. Atoms of one element can combine with atoms of other elements to form _________________. A given compound always has the same relative numbers and types of atoms.

Dalton’s Atomic Theory

5. Atoms are _________________ in chemical processes. – That is, atoms are not created or

destroyed in chemical reactions. – A chemical reaction simply changes the

way the atoms are grouped together.

Dalton’s Atomic Theory

Dalton’s model explained important observations such as the _____________________________.

His model was not accepted at first, however, he used his model to explain the existence of certain types of substances.

He predicted correctly the formation of multiple compounds and his theory became widely accepted.

Atoms

The smallest substance that cannot be divided any further and still maintain the _________________ of the substance.

Structure of an Atom

J.J. (John Joseph) Thomson, physicist– 1890-1900– Showed that the atoms of any element can be made to

emit tiny negative particles - called _________________.

– Thompson knew that the entire atom was not negatively charged so he concluded that the atom must also contain positive particles that balance the negative charge, giving the atom a _________________charge.

J.J. Thomson

– J.J. Thomson was chosen to head the Cavendish Laboratory in Cambridge, England in 1884 when he was only 28 years old.

– Thomson was known for his gift in designing experiments, but he was not mechanically inclined and needed help to build the apparatus needed to perform the experiments.

J.J. Thomson– In the 1890s, one of the most common ways to study

electricity was to build a glass tube with metal electrodes in each end, one of which was coated with zinc sulfide. When the air was pumped out of the tube, called a _________________, and a battery was hooked to wires connected to the electrodes, a bright, glowing spot was observed in the zinc sulfide. After ten years of trying to figure what caused the glow, Thomson finally concluded it must be due to a stream of negative particles that he called _________________- we now call them _________________.

J.J. Thomson

– By the time of Thomson’s discovery in 1897, the Cavendish had become the most distinguished laboratory in England.

– However, despite the prominence of Thomson and his laboratory, the suggestion that these particles came from inside atoms (_________________), was not received by the scientific community - many physicists did not even believe in the existence of atoms.

Structure of an Atom

William Thomson (Lord Kelvin, no relation to J.J. Thomson)– _________________– He had the idea that the atom might be something like a

pudding with raisins randomly distributed throughout. – He reasoned that the atom might be thought of as a

uniform pudding of positives charge with enough negative electrons scattered within to counterbalance that positive charge.

Structure of an Atom

Ernest Rutherford– 1911– Learned physics in J.J. Thomson’s laboratory in the

late 1890s.– Main area of interest was the

_________________ (α particle) - positively charged particles with a mass approximately 7500 times that of an electron.

Ernest Rutherford

– In studying the flight of the α particle through air, Rutherford found that some of the α particles were deflected by something in the air. He designed an experiment that involved directing α particles toward thin metal foil. Surrounding the foil was a detector coated with a substance that produced tiny flashes wherever it was hit by an α particle. The results of the experiment were very different from those he anticipated.

Ernest Rutherford

– Although most of the particles passed straight through the foil, some of the particles were deflected at large angles and some were reflected backwards. (He described this results as comparable to shooting a gun at a piece of paper and having the bullet bounce back!)

– Rutherford knew that if the plum pudding model of the atom was correct, the massive α particles would crash through the thin foil like cannonballs through paper.

Ernest Rutherford

– Rutherford concluded that the plum pudding model for the atom could not be correct.

– The large deflections of the α particles could be caused only by a center of concentrated _________________ charge that would repel the positively charged α particles.

– Most of the α particles passed directly through the foil because the atom is mostly _________________.

Ernest Rutherford– The deflected particles were those that had a “close

encounter” with the positive center of the atom, and the few reflected α particles were those that scored a “direct hit” on the positive center.

– These results could be explained only in terms of a nuclear atom - an atom with a dense center of _________________ (the nucleus) around which tiny electrons moved in a space that was empty.

– He concluded that the nucleus must have a _________________ charge to balance the _________________ charge of the electrons and that it must be small and dense.

Ernest Rutherford

– By 1919, Rutherford concluded that the nucleus of an atom contained what he called _________________ (has the same magnitude of charge as the electron, but its charge is positive)

– _________________ have a 1+ charge and the _________________ a charge of 1-

– 1932, he and a coworker (James Chadwick) were able to show that most nuclei also contain a neutral particle that they named the _________________ (which has no charge)

Modern Concept of Atomic Structure

– The simplest view of the atom is that it consists of a tiny nucleus that is about 10-13 cm in diameter.

– Electrons move about the nucleus at an average distance of about 10-8 cm from it.

– Nucleus contains _________________, which have a positive charge equal in magnitude to the _________________ negative charge, and _________________, which have almost the same mass as a proton but no charge.

Modern Concept of Atomic Structure

Mass and charge of the electron (e-), proton (p+), and neutron (N)

The mass and charge of the electron, proton, and neutron.

Particle Relative Mass* Relative ChargeElectron 1 1-Proton 1836 1+Neutron 1839 None*The electron is assigned a mass of 1 for comparison

Modern Concept of Atomic Structure

If all atoms are composed of these same components, why do different atoms have different chemical properties?– The answer lies in the number and arrangement of the

_________________.– The number of e- a given atom greatly affects the way it can

interact with other _________________.– As a result, atoms of different elements, which have

different numbers of electrons, show different _________________ behavior.

Distinguishing Between Atoms

_________________ and _________________ are equal in an atom of an element (_________________).

The atomic number of an element is the number of _________________ in the nucleus of an atom of that element. (If the p+ and e- are the same, then the atomic number will also identify the number of e-)

Distinguishing Between Atoms

The sum of the number of _________________ and the number of _________________ in a given nucleus is called the atom’s _________________.

Isotopes– atoms with the same number of

_________________ but different numbers of _________________.

– Elements on the periodic table are the most common _________________ of those substances.

Distinguishing Between Atoms

Isotopes– Because they have different numbers of

neutrons, their mass numbers will be different.– Neon - 20– Neon - 21– Neon - 22– All of these are isotopes of neon.

Distinguishing Between Atoms

Isotopes– 3 known isotopes of hydrogen

hydrogen - 1 [hydrogen] hydrogen - 2 [deuterium] hydrogen - 3 [tritium]

Isotopic Symbols

X = the symbol of the element A = the mass number Z = the atomic number

1

1 H

Hydrogen

2

1 H

Deuterium

3

1 H

Tritium

A

Z X

Atomic Masses

Because atoms are so tiny, the normal units of mass - the gram and the kilogram - are much too large to be convenient.

Mass of a single carbon atom is 1.00 x 10-23 grams.

When describing the mass of an atom, scientists have defined a much smaller unit of mass called the _________________- _________________.

Atomic Masses

In terms of grams:– 1 amu = atomic weight of a substance

expressed in grams– 1 carbon atom = 12.01 amu = 12.01 grams– 1 aluminum atom = 26.98 amu = 26.98 grams

Periodic Table of Elements

Shows all the known elements and gives a lot of information about each element.

Invaluable in chemistry!