Atomic Atomic StructureStructure

Chapter 4

Introduction• The view that we have of the atom has

been the result of work done over many years by many people.

• The structure of the atom is composed of smaller subatomic particles.

• Atoms can be distinguished from one another based upon their atomic and mass numbers.

• Isotopes are different versions of a given atom and contribute to the average mass recorded in the periodic table.

Defining the Atom(Section 4.1)

• Early Models of the Atom

• Sizing Up The Atom

The Macro vs. The MicroThe Macro vs. The Micro

What are some differences? What are some similarities?

ElementsElements

• Elements: The simplest form of matter that has a unique set of properties (i.e. different physical and chemical properties). [Section 2.3]

In other words all elements behavedifferently.

Elements = AtomsElements = Atoms

These two terms are used interchangeably.

Therefore, when we talk about an atom of oxygen, we are referring to the element oxygen.

The AtomThe Atom

The smallest particle of an element that retains its identity in a chemical reaction.

The word comes from the Greek word atomos which means “something that cannot be divided further.”

The radii of most atoms fall within the range of 5 x 10-11 m to 2 x 10-10m (50 pm – 200 pm).

STM image of gold atoms

STM image of iron atoms

STM image of platinum atoms

I.) Early Models of the Atom• Democritus at about 400

B.C. proposed that all things were composed of extremely small and irreducible particles. – He lacked experimental

support for his idea and his approach was not based on the scientific method.

These ideas were rejected by Aristotle and were consequently forgotten for almost 2000 years.

• 1803 – John Dalton, an English chemist and school teacher, proposed the modern version of the atomic theory.– Dalton used

experimental methods and transformed Democritus’s ideas on atoms into a scientific theory

Dalton studied the ratios in which elements combine in chemical reactions.

Dalton’s Atomic TheoryDalton’s Atomic Theory1. All elements are composed of tiny

indivisible particles called atoms.

Oxygen atom Hydrogen atoms

2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element. STM Image of Fe

STM image of Au

3. Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds.

4.Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element are never changed into atoms of another element as a result of a chemical reaction.

The electolysis of water

II.) Sizing Up The Atom

Scanning Tunneling Microscope (STM)

How a STM works.

Xenon on Nickel

Carbon Monoxide on Platinum

Is There Anything Is There Anything Smaller Than the Atom?Smaller Than the Atom?

Yes.

These are called subatomic particles.

Structure of the Nuclear Atom(Section 4.2)

• Subatomic Particles

• The Atomic Nucleus

I.) Subatomic ParticlesThe idea of a nuclear atom showed that atoms are divisible.

Subatomic ParticlesSubatomic Particles

Electrons: negatively chargedsubatomic particles.

Protons: positively charged subatomic particles.

Neutrons: subatomic particleswith no charge.

Characteristics of the Subatomic Particles

Discovering the Electron• In 1897 uncovered the existence of electrons using cathode ray tubes (CRT).• Hypothesized that a cathode ray is a stream of negatively charged particles.• He set up an experiment to measure the ratio of the charge of an electron

to its mass. This ratio does not depend

on the kind of gas in the CRT.• Robert Millikan, an American physicist used this ratio to calculate the mass of

the electron.

J.J. ThomsonEnglish Physicist (1856 - 1940)

Thomson's Experiment

CRT with no magnet deflecting stream

CRT with magnet deflecting streamIf electrons are uncharged, then no deflection would occur at all.

II.) The Atomic Nucleus

• Thomson proposed the "plum pudding" model of the atom in which the electrons are stuck into a lump of positive charge.

• No nucleus in this model.

Discovering the Proton

• In 1886, Goldstein found rays travelling in the direction opposite of the cathode rays. He called these rays "canal rays."

• He believed these rays were composed of positive particles. We now know these

particles to be protons.

Eugene GoldsteinGerman Physicist

(1850 - 1930)

• A former student of J.J Thomson.

• Is considered the "father" of nuclear physics.

• Devised the planetary model of the atom, which is Neils Bohr's orbital model.

• Through experiments looking at the scattering pattern of certain particles, he reached his conclusions.

The Rutherford Atom

Ernest Rutherford(1871 - 1937)

• Experiment used alpha particles, which are helium atoms with their electrons removed, thus they have two positive charges.

• The alpha particles are shot at a thin sheet of gold foil (the cannon ball shot at a tissue paper).

• The alpha particles then strike a fluorescent screen behind the gold foil.

Rutherford's Gold-Foil Experiment

Interpreting Rutherford's Results

• If the plum pudding model was true, then almost all the alpha particles would have passed through the foil with only some slight deflection.

• The fact that some particles bounced back indicate that they collided with a similarly charged particle. (The tissue paper sending the cannon ball back at the cannon.

• Conclsion:1. Atom mostly space2. At the center some positive charge exists that

accounts for the deflection. 3. This positively charged region he termed the

nucleus.

• The particles did not all pass through.

• The deflection of the positive alpha particles shows there is a positive charge in the atom.

The Rutherford Model of the Atom

• A centrally located nucleus that contains almost all the mass of the atom and all the positive charge.

• Electrons are negatively charged and distributed around the nucleus. The electrons' orbits accounts for almost all the volume of the atom.

• This model is incomplete and fails to explain the chemical behavior of elements. It needed to be improved upon.

• Worked in Ernest Rutherford's lab before and after WWI.

• Discovered the neutron while studying atomic disintegration. He was looking for a particle with a mass equal to the proton but possessed no charge.

• Werner Heisenberg, using Chadwick's discovery, showed that neutrons were not simply electron/proton pairs, but

their own particles.

Discovering the Neutron

James ChadwickEnglish Physicist

(1891 - 1974)

Label the diagram of the atom

{

Distinguishing Among Atoms(Section 4.3)

• Atomic Number

• Mass Number

• Isotopes

• Atomic Mass

• The Periodic Table – A Preview

I.) Atomic Number & Mass Number

• Elements are different from one another because they contain different numbers of protons.

• Atomic Number: The number of protons in the nucleus of an atom.

• Mass Number: The total number of neutrons and protons in an atom.

Each of these numbers can be found for each element using the periodic table.

The mass number is always the bigger number.

Calculating the number of neutrons in an atom.

• Using the atomic number, the mass number, and the above equation we can determine the number of protons, neutrons, and electrons for any atom.

# of Neutrons = Mass # of Neutrons = Mass number - atomic numbernumber - atomic number

Representing Atomic and Mass Numbers

• Use the symbol for the element

• Place the mass number as a superscript on the left-hand side of the symbol.

• Place the atomic number as the subscript on the left-hand side of the symbol

Au197

79

Gold-197

II.) Isotopes

Atoms that have the same number of protons but different numbers of neutrons.

• Isotopes are chemically similar, but not identical.

• Isotopes of a given element have different mass numbers.

• Each element has a specific number of isotopes (i.e. hydrogen has three naturally occurring isotopes).

• Not all isotopes occur with the same abundance in nature.

Determining the Composition of an Isotope• The composition of an isotope is

determined in the same way as an element.

3 isotopes of chromium exists: chromium-50, chromium-52, and chromium-53. How many neutrons, protons, and electrons are in each isotope, given that the atomic number is 24?

III.) Atomic Masses

• Actual atomic masses are very small – F = 3.155 x 10-23 g. – As = 1.244 x 10-22 g.

• Numbers this small are impractical to work with.

• Better to use relative masses of atoms using carbon-12 as a reference standard. – Carbon-12 = 12 amu.

Atomic Mass Units (amu)

• 1 amu (u) is defined as one-twelfth the mass of a carbon -12 atom.

• 1 amu = 1.66 x 10-24 g.

• Helium = 4.0026 u

Carbon = 12.011 u

If the masses of atoms depend primarily on the number of protons and neutrons, why aren’t the masses listed in the periodic table whole numbers?

In nature most elements occur as a mixture of two or more isotopes, each having a different mass. • Remember, isotopes of an element have different numbers of neutrons and thus will have different masses.

Natural Percent Abundance of Stable Isotopes Natural Percent Abundance of Stable Isotopes of Some Elementsof Some Elements

Name Symbol Natural % Abundance

Mass (amu)

Average atomic mass

Hydrogen 1H

2H

3H

99.985 %

0.015 %

negligible

1.0078 u

2.0141 u

3.0160 u

1.0079

Carbon 12C 13C

98.89 %

1.11 %

12.000 u

13.003 u

12.011

Oxygen 16O17O18O

99.759 %

0.037 %

0.204 %

15.995 u

16.995 u

17.999 u

15.999

Atomic MassesAtomic Masses

• The masses listed in the periodic table are weighted average of the masses of all the elements.

• This weighted average is known as the atomic mass. – A weighted average mass of the atoms in a

naturally occurring sample of an element.

• A weighted mass reflects both the mass and the relative abundance of the isotopes of an element.

Calculating Atomic MassesCalculating Atomic Masses

• To do this calculation you need to know: – the number of stable isotopes

– the mass of each isotope (mi)

– the natural percentage abundance of each isotope (%)

Atomic mass = Σ[(mi) x %]

Let’s practice

Determine the atomic mass of carbon. Carbon has two stable isotopes.

Carbon-12 = 12.000 u (98.99%)

Carbon-13 = 13.003 u (1.11%)

IV.) The Periodic Table – A Preview• The periodic table is an arrangement of

elements in which the elements are separated into groups based on a set of repeating properties.

• A periodic table allows us to easily compare the properties of one element (or group of elements) to another element (or group of elements).

We will discuss the periodic table in greater detail in chapter 6.

(7 periods)(18 periods)

The periodic table gives a wealth of info. regarding the elements. Such as:

• Atomic and mass numbers.

• Chemical properties of each element.

• The sizes of the atoms.

• The location of electrons in each atom.

Use the periodic table to locate the following elements, given their atomic numbers.

1. Atomic Number = 29

2. Atomic Number = 50

3. Atomic Number = 82

4. Atomic Number = 79

Now write their chemical symbols. Include the atomic number and mass number.

Atomic Atomic StructureStructure

Chapter 4

The End