Atoms Chapter 4

1. Atoms Chapter 4 Section 1—The Development of Atomic Theory Section 2—The Structure of Atoms Section 3—Modern Atomic Theory

2. The Development of Atomic Theory • Atoms are everywhere. They make up the air we breath and the chairs we sit in, and the clothes we wear. Atoms determine the properties of matter.

3. The Beginnings of Atomic Theory • The atomic theory was developed slowly over a long period of time. The first theory of atoms was proposed over 2,000 years ago. • In the fourth century BCE, the Greek philosopher Democritus suggested that the universe was made of individual units. • He called these units atoms. “Atom” comes from atomos, a Greek word that means “unable to be cut or divided”.

4. Democritus did not have evidence for his atomic theory. • His theory of atoms explained some observations, he did not have the evidence needed to convince people that atoms existed. • Some people supported his theory and other theories were also proposed

5. As the science of chemistry is developing in the 1700’s the emphasis on making careful and repeated measurements in experiments increased. • Because of this change, more precise data is collected and it was used to favor one theory over another.

6. Dalton’s Atomic Theory • In 1808 an English schoolteacher named John Dalton proposed a revised theory. His theory was developed on scientific basis, and some parts of his theory hold true today. • Like Democritus, Dalton proposed that atoms could not be divided. • According to Dalton, all atoms of a given element were exactly alike, and atoms of different elements could join to form compounds.

7. How are Dalton’s and Democritus atomic theories similar? • Both believed that atoms are the fundamental units of matter and that atoms are indivisible.

8. Dalton used experimental evidence • Unlike Democritus, Dalton based his theory on experimental evidence. Scientist were beginning to observe that some substances combined together in consistent ways. • According to the law of definite proportions, a chemical compound always contains the same elements in exactly the same proportions by weight or mass. Example is H₂O or water it contains the same proportions of hydrogen and oxygen by mass. This and other evidence supported Dalton’s theory.

9. Daltons theory did not fit all the observations • His theory is considered the foundation for modern atomic theory. • Some parts of this theory were correct, but could not explain all the experimental evidence. • Like many scientific theories, the atomic theory changed over the years as new experiments were performed and acquired new information.

10. Thompson’s Model of the Atom • In 1897 J. J. Thompson, a British scientist conducted an experiment that suggested that atoms were not indivisible. • He was experimenting with electricity and was not planning on learning about atoms. • He was studying cathode rays, mysterious rays in a vacuum tube.

11. Thompson’s cathode-ray experiment suggested that cathode rays were made of negatively charged particles that came from inside atoms. • This result revealed that atoms could be divided into smaller parts.

12. Thompson developed the plum-pudding model • Two metal plates at the ends of a vacuum tube are called the cathode and anode. The cathode has a negative charge and the anode a positive charge. When voltage is applied across the plates, a glowing beam comes from the cathode and strikes the anode. • Thompson knew that magnets deflect charges. He reasoned that because the air was removed from the tube, the beam must have come from the cathode or from the anode.

13. The direction of the deflection confirmed that the beam was made of negative charges and therefore came from cathode. • Thompson had discovered the electrons, negatively charged particles inside the atom. Electron is a subatomic particle with a negative charge. • His model was called the plum-pudding model after a dessert popular in his day.

14. Rutherford’s Model of Atoms • Soon after Thompson proposed his new atomic model, Ernest Rutherford, another British scientist, developed an experiment to test Thompson’s model. He found that Thompson’s model needed to be revised. • Rutherford proposed that most of the mass of the atom was concentrated at the atom’s center.

15. Rutherford conducted the gold-foil experiment • Two of Rutherford’s students aimed a beam of positively charged alpha particles at a very thin sheet of gold foil. He predicted that most of the particles would travel in a straight path and a few would be slightly deflected.

16. The observations from the experiment did not match Rutherford’s predictions. • Most particles passed through the gold foil, but some were deflected by a large amount. A few particles came straight back. • Why were Rutherford’s results so surprising? • Because they did not match his predictions based on Thompson’s model of the atom. He expected most of the positive particles to pass straight through, but instead, several were deflected at large angles.

17. What do you think he discovers? • Think about those positively charged particles being aimed at the gold fold. • The nucleus • This is the atoms central region, which is made up of protons and neutrons.

18. Rutherford discovers the nucleus. • His experiment suggested that an atom’s positive charge is concentrated in the center of the atom. • This positively charged, dense core is called the nucleus. • Incoming positive charges that passed close to the nucleus were deflected sharply. Incoming positive charges that were aimed directly at the nucleus bounced straight back. • Why? • Because same charges repel each other.

19. The data from Rutherford’s experiment suggested that the nucleus was small compared to the rest of the atom. • In Rutherford’s results led to a new model of the atom. His model had negative electrons orbiting the positively charged nucleus, like planets orbiting the sun. • Today we know that the nucleus has protons and neutrons. Protons have a positive charge and neutrons have no charge.

20. The Structure of AtomsSection 2 • Less than 100 years after Dalton published his atomic theory, scientist determine that atoms consist of smaller particles, such as electrons. • Why it matters—Radioisotopes emit energy when they decay. To diagnose and treat diseases, doctors use this property of radioisotopes to track where in the body certain atoms go.

21. What is in an Atom • Atoms are made up of various subatomic particles. • The three main subatomic particles are distinguished by mass, charge, and location in the atom. • At the center of each atom is a small, dense nucleus. • The nucleus is made up of protons and neutrons. • Proton is a subatomic particle that has a positive charge and that is located in the nucleus of an atom. • Neutron is a subatomic particle that has no charge and that is located in the nucleus of an atom.

22. Protons and neutrons are almost identical in size and mass. Moving around outside the nucleus is a cloud of very tiny negatively charged electrons. • The mass of an electron is much smaller than that of a proton or neutron.

23. Each element has a unique number of protons and is defined by the number of protons in an atom of that element. Example: Carbon has 6 protons. • Unreacted atoms have no overall charge. Most atoms do not have an overall charge. The reason is that most atoms have an equal number of protons and electrons so these charges exactly cancel.

24. Example helium atom: it has 2 protons and 2 electrons and they cancel each other out for no charge. • Charge of two protons +2 • Charge of two neutrons 0 • Charge of two electrons -2 • Total charge of a helium 0 atom. • If an atom gains or loses electrons, it becomes charged. A charged atom is called an ion.

25. The electric force holds the atom together. • The positive and negative charges attract each other with a force known as the electric force. • The electric force between protons and electrons holds the atom together. • On a larger scale this same force holds solid and liquid materials together. • The electric attraction hold water molecules together.

26. Atomic Number and Mass Number • Atoms of different elements have their own unique structure. • Atoms of each element have the same number of protons, but they can have different numbers of neutrons.

27. The atomic number is the number of protons in the nucleus of an atom. • The atomic number also equals the number of electrons. • Each element is defined by its unique number of protons each element has a unique atomic number. • The largest naturally occurring element is uranium with 92 protons.

28. The mass number equals the total number of subatomic particles in the nucleus. • Mass number is the sum of the numbers of protons and neutrons in the nucleus of an atom.

29. The mass number reflects the number of protons and neutrons because these provide most of the atoms mass. Electrons are not included. • Atoms of an element have the same atomic number they can have a different mass number because the number of neutrons can vary.

30. Which defines an element: the atomic number of an element or the mass number of an element? • The atomic number defines the element because atoms of each element always have the same number of protons but can have different numbers of neutrons.

31. Isotopes • An isotope is an atom that has the same number of protons (or the same atomic number) as other atoms of the same element do but that has a different number of neutrons (and thus a different atomic mass). • Because they have the same number of protons, isotopes of an element have the same chemical properties.

32. Isotopes of an element vary in mass because their number of neutrons differ. • Each isotope of hydrogen has one proton, but the number of neutrons varies. Isotopes of hydrogen Protium –has 1 proton, 1 electron, and no neutrons. Has mass number of 1. Deuterium—has 1 neutron, 1 proton, an 1 electron. Has mass number of 2. Tritium—has 2 neutrons, 1 proton, and 1 electron. Has a mass number of 3.

33. Some isotopes are more common than others. • In hydrogen the protium is the most common isotope both on the sun and on Earth. • 1 out of 6,000 in Earth’s crust is a deuterium isotope and tritium is unstable that decays over time. It is called a radioisotope, it emits radiation and decay into other isotopes. It will continue to decay until it reaches a stable form. • Each radioisotope decays at a fixed rate which can vary from a fraction of a second to millions of years.

34. The number of neutrons can be calculated. • If you know the atomic number and the mass number you can calculate the number of neutrons an atom has. • You take the mass number and subtract the atomic number and this leaves you with the number of neutrons. • Let’s look in out book on page 124.

35. Atomic Masses • The mass of a single atom is very small. • Because working with such tiny masses is difficult atomic masses are usually expressed in unified atomic mass units. • It is equal to one-twelfth of the mass of a carbon-12 atom. This is called the atomic mass unit, amu. • Unified atomic mass unit is a unit of mass that describes the mass of an atom or molecule; it is exactly 1/12 the mass of a carbon atom with a mass number of 12. (symbol, u).

36. Average atomic mass is a weighted average. • Often the atomic mass listed for an element on the periodic table is an average atomic mass for the element as found in nature. The average atomic mass for an element is a weighted average. • Which means commonly found isotopes have a greater effect on the average atomic mass that rarely found isotopes have a greater effect on the average atomic mass that rarely found isotopes do.

37. The mole is useful for counting small particles. • Because chemists often deal with large numbers of small particles, they use a large counting unit called the mole (mol). A mole is a collection of a very large number of particles. • It equals 602,213,679,000,000,000,000,000 particles. • But written as 6.022 x 10²³ this is called Avogadro’s number. • The mole has been defined as the number of atoms in 12.00 grams of carbon-12. so one mole contains 6.022 x 10²³ particles of that substance.

38. Avogadro’s number is useful for counting atoms or molecules. • How many particles are in 1 mol of iron? • A mole of iron contains 6.022 x 10²³ iron atoms. • Mole--is the SI base unit used to measure the amount of a substance whose number of particles is the same as the number of atoms of carbon in exactly 12g of carbon-12.

39. Moles and grams are related. • The mass in grams of one mole of a substance is called molar mass. • In nature, elements often occur as mixtures of isotopes. So, a mole of an element usually contains several isotopes. As a result, an elements molar mass in grams per mole equals its average atomic mass in unified atomic mass units u.

40. You can convert between moles and grams. • See your book page 126, figure 8 which shows you how to make the conversion. • There are three steps to converting grams of a substance to moles. • Determine how many grams are given in the problem. • Calculate the molar mass of the substance. • Divide step one by step two. • The three steps above can be expressed in the following proportion:

41. Grams of the substance Molar mass of a substance in grams. ______________________ = ______________________ Moles of the substance One mole See text book page 126 for practice problems.

42. Compounds also have molar masses. • To find the molar mass of a compound you can add up the molar masses of all the atoms in a molecule compound. • See water example in text book: page 127

43. Modern Atomic Theory Section 3 • Like most scientific models and theories the model of the atom has been revised many times to explain new discoveries.

44. Modern Models of the Atom • In the modern atomic model, electrons can be found only in certain energy levels. Furthermore, the locations of electron cannot be predicted precisely.

45. Electron location is limited to energy levels. • In 1913, Niels Bohr, a Danish physicist, suggested that the energy of each electron was related to the electron’s path around the nucleus. Electrons can be in only certain energy levels. They must gain energy to move to a higher energy level or must lose energy to move to a higher energy level or must lose energy to move to a lower level. • His description of energy levels is still used by scientist today.

46. Electrons act like waves. • A new model no longer assumes that electrons no longer orbit the nucleus in a definite path in the same way planet orbit the sun, was proposed. • The new model shows electrons behave more like waves on a vibrating string than like particles.

47. How does the electron wave model of the atom differ from earlier atomic models? • In earlier atomic models, electrons were considered to be particles. In the electrons-wave model, electrons act more like waves than like particles.

48. The exact location of electrons cannot be determined. • The exact location of electrons in an atom and the speed and direction is impossible to locate. • Best that a scientist can do is calculate the chance of finding an electron in a certain place within an atom. • Orbital--is a region in an atom where there is a high probability of finding electrons.

49. Electron Energy Levels • Within an atom, electrons have various amounts of energy exist in different energy levels. • There are many possible energy levels that an electron can occupy. • The number of energy levels that are filled in an atom depends on the number of electrons. • The electrons in the outer energy level of an atom are called valance electrons.

50. Valence electrons determine the chemical properties of an atom. • Valence electrons--an electron that is found in the outer most shell of an atom and that determines the atom’s chemical properties.