Chemistry Unit I: Chapter 4 “Atoms”

1. Chemistry Unit I:Chapter 4 “Atoms”

2. Chp 4.1 Atomic Structure, p. 113-118 • Our understanding of atoms took many centuries. • The 4th century Greek philosopher, Democritus, suggested that the universe was made of invisible units called “atoms”. The word atom means “that which cannot be divided”. • The Greek idea of atoms had to wait 2000 years before it became accepted.

3. John Dalton • In 1808, an English schoolteacher named John Dalton proposed his own atomic theory. He used evidence such as the “law of definite proportions”. • His theory proposed: • Every ELEMENT is made of tiny, unique particles called atoms that cannot be subdivided. • Atoms of the same element are exactly alike. • Atoms of different elements can join to form molecules.

4. J.J. Thomson • In 1897, British scientist, J.J. Thomson, found that atoms contain negatively charged particles. These particles were named “electrons”. • Since atoms are neutral, Thomson reasoned that atoms must contain some sort of positive charge. (Cathode Ray Tube experiment) • His model, like a plum pudding, had negative electrons (plums) embedded in a positive sphere (pudding).

5. Ernest Rutherford • In 1911, Ernest Rutherford, discovered that atoms contain a dense, positively charged “nucleus” through his Gold Foil Experiment. • Later Rutherford name the positive particles “protons”. • In his model, atoms were mostly empty space with electrons moving around a small, very dense, positively-charged nucleus in the center of the atom.

6. One addition…the neutron • In 1932, British scientist, James Chadwick, discovered another particle in the nucleus of an atom. The particle was nearly the same mass as a proton. • This particle is called a “neutron”, because it is electrically neutral.

7. Chapter 4.2 “The Structure of Atoms,” p.119-127 • Three main subatomic particles are distinguished by mass, charge and location in the atom. • Proton, charge +1, located in the nucleus, mass = 1.67x10-27 • Neutron, charge 0, located in the nucleus, mass= 1.67x10-27 • Electron, charge -1, located in orbitals, mass = 9.11x10-31

8. Atoms • Each element has a unique number of protons. • Each atom of the element will have the same number of protons. • Atoms are electrically “neutral” overall. Atoms have the same number of protons as electrons. These charges exactly cancel leaving no charge. • Positive and negative charges attract each other with a force known as “electric force”.

9. Ions • Atoms may undergo a process called “ionization”, in which they lose or gain electrons in order to fill their valence energy level. • An “ion” is an atom (or group of atoms) that has lost or gained one or more electrons, therefore has a net charge. • Removing electrons, causes a positive ion to form, aka cation. • Adding electrons, causes a negative ion to form, aka anion.

10. Atomic Nucleus • The atomic number tells you how many protons are in an atom. • Every element has a different atomic number. • Atoms may or may not have the same number of neutrons as protons. • The mass number tells you how many protons plus neutrons are in the nucleus of an atom all together.

11. Isotopes • Some atoms contain a different number of neutrons thereby making them “heavier” or “lighter” than the average atom. These atoms are called isotopes. • Isotopes have the same number of protons just a different number of neutrons. • To calculate the number of neutrons, subtract the atomic number (Z) from the mass number (A) like this: • A – Z = number of neutrons

12. Atomic Masses • The mass of a single atom is extremely small! • Atomic masses are usually expressed in unified atomic mass units. • A unified atomic mass unit (u) is equal to one-twelfth of the mass of a carbon-12 atom. Aka “amu” • The average atomic mass for an element is a weighted average and appears as a decimal number on the periodic table.

13. Moles • Scientists often deal with very large numbers of small particles. They use a counting unit called a “mole”. • A mole is 602213670000000000000000 !!! • This number is usually written as 6.02 x 10 23 /mole and is called Avogadro’s constant. This is how many particles are in a mole of any substance. • Like the word “dozen” means 12, a mole of a substance contains 6.02 x 10 23 particles.

14. Molar Mass • The mass in grams of 1 mol of a substance is called its molar mass. • Example, one mole of lithium has a mass of 6.941 g. • Because mass and amount of a substance are related, it is possible to convert moles to grams and vice versa.

15. Calculating with Moles • Using a conversion factor allows you to change from mass to amount and back again. • The numerator of the conversion factor contains the unit you want to change to. The denominator of the conversion factor contains the unit you start with. • The numerator and denominator equal 1 since they are equal. • Example: you could use either • 10 gumballs/ 21.4 g • 21.4 g/ 10 gumballs

16. Calculating with moles • Converting between amount of an element in moles and its mass in grams you can use this chart: Amount (mol) x molar mass (g)/ 1 mol of element = Mass (g) Mass (g) x 1 mol of element/ molar mass of element (g) = Amount (mol)

17. Chp 4-3 Modern Atomic Theory p.128-132 • Modern Model of the Atom • Electrons can be found only in certain energy levels, not between levels. • Furthermore, the location of electrons cannot be predicted precisely.

18. Niels Bohr • In 1913, Niels Bohr, a Danish scientist, revised the Rutherford’s model. • Bohr showed that electrons could only have specific amounts of energy, leading them to move in certain orbits. • This model resembled planets orbiting the sun.

19. Cloud of Electrons • In the 1920’s, scientists again revised the model. They found that electrons can be anywhere in a cloudlike region around the nucleus. This region is called an orbital. • Electrons behave more like waves on a vibrating string than like particles. • The different energy levels contain different numbers of electrons depending on the energy level size and electron energy.

20. Electron Energy Levels • The number of energy levels that are filled in an atom depends on the number of electrons. • The electron(s) in the outermost energy level of an atom are called “valence electrons”. • The most valence electrons that an atom can have is eight (8). • Atoms with a filled valence energy level are the “stable” and don’t react chemically. • Electron dot diagrams show the element’s symbol and valence electrons as dots.

21. Electrons and Orbitals • Electrons may occupy four different kinds of orbitals within atoms: • “S” is the simplest orbital, a sphere shape, lowest energy. • “p” is dumbbell-shaped, and can be oriented 3 ways, and are the next higher energy levels. • “d” and “f” orbitals are much more complex. There are 5 possible “d” orbitals and 7 possible “f” orbitals. These are the highest energy levels. • These orbitals are all different shapes and can hold a maximum of two electrons each.

22. Electron Transition • Electrons jump between energy level when an atom gains or loses energy. • Ground state is the lowest state of energy of an electron. • If an electron gains energy, it moves to an excited state. • It gains energy by absorbing a particle of light called a photon. • It will release the photon when it moves back to its ground state.

23. Stop for Now!!!! Chapter 4 content complete • (the following slides are from Chp 5…please skip for now)

24. Chapter 5.3 “Families of Elements,” p 156-165 • The two main groups on the periodic table are metals and non-metals. • Most metals are shiny solids that can be stretched (ductile) and shaped (malleable). The are good conductors of heat and electricity. • Non-metals can be solid, liquid or gas. They are typically dull, brittle and are poor conductors. • Between the two groups are semiconductors, or metalloids, which share characteristics of both groups.

25. Alkali Metals • Group 1 of the periodic table contains elements with one valence electron, which can be easily removed to form a positive ion. • They are all highly reactive and are not found in nature as elements. • Examples include sodium and potassium. These easily combine with non-metals to form “salts”.

26. Alkaline Earth Metals • Group 2 of the Periodic Table are the alkaline earth metals. • They contain 2 valence electrons and react (not as violently as group 1) losing their 2 electrons to become a positive ion then forming compounds. • Examples in calcium and magnesium.

27. Transition Metals • Groups 3 through 12 contain the transition metals. • They are less reactive than groups 1 or 2 but some still combine with non-metals to form compounds. Some have many versions of cations. • Examples include gold, silver, copper mercury etc.

28. Non-metals • Except for hydrogen, non-metals are found on the right side of the periodic table including some elements in groups 13 thru 16 and all elements in groups 17 and 18. • Non-metals tend to share or gain electrons (anion) and are plentiful on Earth. • Examples include hydrogen, carbon, oxygen, sulfur, nitrogen, etc.

29. Halogens • Group 17 contain non-metals called halogens. These elements contain 7 valence electrons and are highly reactive. • Examples include fluorine, chlorine, iodine, etc.

30. Noble Gases • Group 18 of the periodic table contain the noble gases. Because these atoms have filled valence energy levels, they do not react with other elements to form compounds….hence the name “noble”. • They all exist as single gas atoms. • Examples include helium, neon, argon, etc.

31. Semiconductors • Located along the “staircase line”, semiconductors or metalloids have properties of both metals and non-metals. • Semiconductors are able to conduct heat and electricity under certain conditions, some are hard, some are shiny and so on. • The 6 elements are: • Boron, silicon, germanium, arsenic, antimony, and tellurium. Sometimes polonium and astatine are also included.