John Dalton (1766- 1844), an English schoolteacher and chemist, studied the results of experiments...

• John Dalton (1766-1844), an English schoolteacher and chemist, studied the results of experiments by many other scientists.

Dalton’s Atomic Theory

Atomic Structure: Basic ConceptsAtomic Structure: Basic Concepts Atomic Structure: Basic ConceptsAtomic Structure: Basic Concepts Topic 2Topic 2

• Dalton proposed his atomic theory of matter in 1803.

• Although his theory has been modified slightly to accommodate new discoveries, Dalton’s theory was so insightful that it has remained essentially intact up to the present time.



• The following statements are the main points of Dalton’s atomic theory.


1. All matter is made up of atoms.

2. Atoms are indestructible and cannot be divided into smaller particles. (Atoms are indivisible.)

3. All atoms of one element are exactly alike, but are different from atoms of other elements.


• 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.

The Electron

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

• Thomson’s experiments used a vacuum tube.


• A vacuum tube has had all gases pumped out of it.

The Electron

• At each end of the tube is a metal piece called an electrode, which is connected through the glass to a metal terminal outside the tube.

• These electrodes become electrically charged when they are connected to a high-voltage electrical source.


• When the electrodes are charged, rays travel in the tube from the negative electrode, which is the cathode, to the positive electrode, the anode.

Cathode-Ray Tube

• Because these rays originate at the cathode, they are called cathode rays.


• Thomson found that the rays bent toward a positively charged plate and away from a negatively charged plate.

Cathode-Ray Tube

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


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• These electrons had to come from the matter (atoms) of the negative electrode.

Cathode-Ray Tube

• Thomson concluded that cathode rays are made up of invisible, negatively charged particles referred to as electrons.


• From Thomson’s experiments, scientists had to conclude that atoms were not just neutral spheres, but somehow were composed of electrically charged particles.

Cathode-Ray Tube

• Reason should tell you that there must be a lot more to the atom than electrons.

• Matter is not negatively charged, so atoms can’t be negatively charged either.


• If atoms contained extremely light, negatively charged particles, then they must also contain positively charged particles—probably with a much greater mass than electrons.

Cathode-Ray Tube


• In 1886, scientists discovered that a cathode-ray tube emitted rays not only from the cathode but also from the positively charged anode.

Protons

• These rays travel in a direction opposite to that of cathode rays.


• Like cathode rays, they are deflected by electrical and magnetic fields, but in directions opposite to the way cathode rays are deflected.

Protons

• Thomson was able to show that these rays had a positive electrical charge.

• Years later, scientists determined that the rays were composed of positively charged subatomic particles called protons.


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

Protons

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


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 experimenters set up a lead-shielded box containing radioactive polonium, which emitted a beam of positively charged subatomic particles through a small hole.



• The sheet of gold foil was surrounded by a screen coated with zinc sulfide, which glows when struck by the positively charged particles of the beam.

• Today, we know that the particles of the beam consisted of clusters containing two protons and two neutrons and are called alpha particles.



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.


The Nuclear Model of the Atom• The new model of the atom as pictured by

Rutherford’s group in 1911 is shown below.


Atomic Numbers

• It is the number of protons that determines the identity of an element, as well as many of its chemical and physical properties.

• The atomic number of an element is the number of protons in the nucleus of an atom of that element.


Atomic Numbers

• Therefore, the atomic number of an element also tells the number of electrons in a neutral atom of that element.

• Because atoms have no overall electrical charge, an atom must have as many electrons as there are protons in its nucleus.


• The sum of the protons and neutrons in the nucleus is the mass number of that particular atom.

Masses

• The mass of a neutron is almost the same as the mass of a proton.


Atomic Mass

• In order to have a simpler way of comparing the masses of individual atoms, chemists have devised a different unit of mass called an atomic mass unit, which is given the symbol u.

• An atom of the carbon-12 isotope contains six protons and six neutrons and has a mass number of 12.


Atomic Mass

• Chemists have defined the carbon-12 atom as having a mass of 12 atomic mass units.

• Therefore, 1 u = 1/12 the mass of a carbon-12 atom.

• 1 u is approximately the mass of a single proton or neutron.


Information in the Periodic Table

• The number at the bottom of each box is the average atomic mass of that element.

• This number is the weighted average mass of all the naturally occurring isotopes of that element.


Electrons in Motion

• Niels Bohr (1885-1962), a Danish scientist who worked with Rutherford, proposed that electrons must have enough energy to keep them in constant motion around the nucleus.

• Electrons have energy of motion that enables them to overcome the attraction of the positive nucleus.


Electrons in Motion

• Bohr’s view of the atom, which he proposed in 1913, was called the planetary model.

• This energy keeps the electrons moving around the nucleus.


The Electromagnetic Spectrum

• To boost a satellite into a higher orbit requires energy from a rocket motor.

• One way to increase the energy of an electron is to supply energy in the form of high-voltage electricity.

• Another way is to supply electromagnetic radiation, also called radiant energy.



• Radiant energy travels in the form of waves that have both electrical and magnetic properties.

• These electromagnetic waves can travel through empty space, as you know from the fact that radiant energy from the sun travels to Earth every day.



• As you may already have guessed, electromagnetic waves travel through space at the speed of light, which is approximately 300 million meters per second.



• Electromagnetic radiation includes radio waves that carry broadcasts to your radio and TV, microwave radiation used to heat food in a microwave oven, radiant heat used to toast bread, and the most familiar form, visible light.

• All of these forms of radiant energy are parts of a whole range of electromagnetic radiation called the electromagnetic spectrum.


Electrons and Light• The spectrum of light released from excited

atoms of an element is called the emission spectrum of that element.


Evidence for Energy Levels

• Bohr theorized that electrons absorbed energy and moved to higher energy states.

• Then, these excited electrons gave off that energy as light waves when they fell back to a lower energy state.


Evidence for Energy Levels

• These regions of space in which electrons can move about the nucleus of an atom are called energy levels.

• Because electrons can have only certain amounts of energy, Bohr reasoned, they can move around the nucleus only at distances that correspond to those amounts of energy.


The Electron Cloud Model

• Instead, they are spherical regions of space around the nucleus in which electrons are most likely to be found.

• As a result of continuing research throughout the 20th century, scientists today realize that energy levels are not neat, planetlike orbits around the nucleus of an atom.


• The space around the nucleus of an atom where the atom’s electrons are found is called the electron cloud.


• These spherical regions where electrons travel may be depicted as clouds around the nucleus.

• Electrons themselves take up little space but travel rapidly through the space surrounding the nucleus.


Electrons in Energy Level

• Each energy level can hold a limited number of electrons.

• How are electrons arranged in energy levels?

• The lowest energy level is the smallest and the closest to the nucleus.



• The second energy level is larger because it is farther away from the nucleus. It holds a maximum of eight electrons.

• This first energy level holds a maximum of two electrons.

• The third energy level is larger still and holds a maximum of 18 electrons.


Energy Levels• A hydrogen atom has only one electron. It’s

in the first energy level.



• You can also use the periodic table as a tool to predict the number of valence electrons in any atom in Groups 1, 2, 13, 14, 15, 16, 17, and 18.

• The electrons in the outermost energy level are called valence electrons.

• All atoms in Group 1, like hydrogen, have one valence electron. Likewise, atoms in Group 2 have two valence electrons.


Electrons in Energy Level• An oxygen atom has eight electrons. Two of

these fill the first energy level, and the remaining six are in the second energy level.


Lewis Dot Diagrams

• Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols.


Lewis Dot Diagrams

• A Lewis dot diagram illustrates valence electrons as dots (or other small symbols) around the chemical symbol of an element.


Lewis Dot Diagrams• Each dot represents one valence electron. • In the dot diagram, the element’s symbol

represents the core of the atom—the nucleus plus all the inner electrons.


Question 1

How does the atomic number of an element differ from the element’s mass number?

Answer

The atomic number of an element is the number of protons in the nucleus. The mass number is the sum of the number of protons and neutrons.

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Basic Concept QuestionsBasic Concept Questions

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A. Chlorine

Write a Lewis dot diagram for each of the following.

Question 2

C. Potassium

B. Calcium

A. Chlorine

Answer

C. Potassium

B. Calcium

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A. Ultraviolet light

Give an example for each type of electromagnetic energy listed below.

Question 3

C. Visible light

B. Infrared light

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Sample answers:

A. ultraviolet light:

Answer

the spectrum of lightwe see as color

C. visible light:

B. infrared light:

part of sunlight

radiant heat

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Atomic Structure: Additional ConceptsAtomic Structure: Additional ConceptsAtomic Structure: Additional ConceptsAtomic Structure: Additional Concepts

Additional Concepts

Energy Levels and SublevelsTopic 2Topic 2


• The emission spectrum for each element has a characteristic set of spectral lines.

• This means that the energy levels within the atom must also be characteristic of each element.

• But when scientists investigated multi-electron atoms, they found that their spectra were far more complex than would be anticipated by the simple set of energy levels predicted for hydrogen.

Energy Levels and Sublevels

• Notice that these spectra have many more lines than the spectrum of hydrogen.

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• Some lines are grouped close together, and there are big gaps between these groups of lines.

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• The big gaps correspond to the energy released when an electron jumps from one energy level to another.

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• This suggests that sublevels—divisions within a level—exist within a given energy level.


• The interpretation of the closely spaced lines is that they represent the movement of electrons from levels that are not very different in energy.

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• If electrons are distributed over one or more sublevels within an energy level, then these electrons would have only slightly different energies.

• The energy sublevels are designated as s, p, d, or f.

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• Each energy level has a specific number of sublevels, which is the same as the number of the energy level.

• For example, the first energy level has one sublevel. It’s called the 1s sublevel.

• The second energy level has two sublevels, the 2s and 2p sublevels

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• Within a given energy level, the energies of the sublevels, from lowest to highest, are s, p, d, and f.

• The third energy level has three sublevels: the 3s, 3p, and 3d sublevels; and the fourth energy level has four sublevels: the 4s, 4p, 4d, and 4f sublevels.

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The Distribution of Electrons in Energy Levels

• A specific number of electrons can go into each sublevel.

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• An s sublevel can have a maximum of two electrons, a p sublevel can have six electrons,

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• a d sublevel can have ten electrons, and an f sublevel can have 14 electrons.

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Orbitals

• This principle is known as the Heisenberg uncertainty principle. In 1932, Heisenberg was awarded the Nobel Prize in Physics for this discovery, which led to the development of the electron cloud model to describe electrons in atoms.

• In the 1920s, Werner Heisenberg reached the conclusion that it’s impossible to measure accurately both the position and energy of an electron at the same time.

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OrbitalsTopic 2Topic 2


• The electron cloud model is based on the probability of finding an electron in a certain region of space at any given instant.

• In any atom, electrons are distributed into sublevels and orbitals in the way that creates the most stable arrangement; that is, the one with lowest energy.

Electron Configurations• This most stable arrangement of electrons in

sublevels and orbitals is called an electron configuration.

• Electrons fill orbitals and sublevels in an orderly fashion beginning with the innermost sublevels and continuing to the outermost.

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Building Electron Configurations

• The electron configuration for carbon is 1s22s22p2.

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Building Electron ConfigurationsTopic 2Topic 2


Building Electron Configurations

• At element number 10, neon, the p sublevel is filled with six electrons.

• The electron configuration for neon is 1s22s22p6.

• Neon has eight valence electrons; two are in an s orbital and six are in p orbitals.

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Noble Gases• Each period

ends with a noble gas, so all the noble gases have filled energy levels and, therefore, stable electron configurations.

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Noble Gases

• These stable electron configurations explain the lack of reactivity of the noble gases. Noble gases don’t need to form chemical bonds to acquire stability.

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Calculating Atomic MassTopic 2Topic 2


Calculating Atomic Mass

• Copper exists as a mixture of two isotopes.

• The lighter isotope (Cu-63), with 29 protons and 34 neutrons, makes up 69.17% of copper atoms.

• The heavier isotope (Cu-65), with 29 protons and 36 neutrons, constitutes the remaining 30.83% of copper atoms.

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• The atomic mass of Cu-63 is 62.930 amu, and the atomic mass of Cu-65 is 64.928 amu.

• Use the data above to compute the atomic mass of copper.

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Calculating Atomic Mass• First, calculate the contribution of each

isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

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• The average atomic mass of the element is the sum of the mass contributions of each isotope.

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Additional Assessment QuestionsAdditional Assessment Questions

Write electron configurations and abbreviated electron configurations of the following elements.

Question 1

A. Boron

C. Phosphorus

B. Fluorine

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Answer

A. Boron

C. Phosphorus

B. Fluorine

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Question 2

The table on the next slide shows the five isotopes of germanium found in nature, the abundance of each isotope, and the atomic mass of each isotope.

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Calculate the atomic mass of germanium.

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Answer

72.59 amu

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John Dalton (1766- 1844), an English schoolteacher and chemist, studied the results of experiments...

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