1 UNIT 2 Atoms, Molecules, and Ions

2 The Power of 10 nceopticsu/powersof10/

3 What to study for Unit 2 Look at syllabus (Objectives) List of ions to learn Homework recommendations Online Quiz, due Monday, 9/15 (5 pm) Exam Tuesday, 9/16 Don’t forget Unit 1!!!

4 A HISTORY OF THE STRUCTURE OF THE ATOM

5 History Greek Philosopher Democritus ( B.C.): all matter composed of small atoms atomos = indivisible 1803, John Dalton (brit.): atoms are the fundamental building blocks of matter

6 Dalton's Postulates Each element is composed of extremely small particles called atoms.

7 Dalton's Postulates All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

8 Dalton's Postulates Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

9 Dalton’s Postulates Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

10 John Dalton’s Atomic Theory (ca 1803) 1.Each element is composed of extremely small particles called atoms. 2.All atoms of a given element are identical. The atoms of different elements are different and have different properties (including different masses). 3.Atoms of an element are not changed into different types of atoms by chemical reactions. Atoms are neither created nor destroyed in chemical reactions. This is the Law of Conservation of Mass. 4.Compounds are formed when atoms of more than one element combine. A given compound always has the same relative number and kind of atoms. This is the Law of Constant Composition.

11 John Dalton’s Atomic Theory Led him to deduce the Law of Multiple Proportions: When two or more elements combine to form more than one compound, the relative masses of the elements which combine will be in in the ratio of small whole numbers. In carbon monoxide, CO, 12 g carbon combine with 16 g oxygen. C:O ratio is 12:16 or 3:4. In carbon dioxide, CO 2, 12 g carbon combine with 32 g oxygen. C:O ratio is 12:32 or 3:8.

12 John Dalton’s Atomic Theory Almost right. A good start. Structure of the atom after Dalton (ca. 1810) very small

13 J.J. Thomson (1897): Cathode Rays Atoms subjected to high voltages give off cathode rays.

14 J.J. Thomson: Cathode Rays Cathode rays can be deflected by a magnetic field. Cathode rays are negatively charged particles (electrons). Electrons are in atoms.

15 J.J. Thomson – The Electron Structure of the atom after Thomson (ca. 1900) “Plum pudding” model: Negative electrons are embedded in a positively charged mass. Positively charged mass Electrons (-) Unlike electrical charges attract, and that is what holds the atom together.

16 Radioactivity Radioactivity is the spontaneous emission of radiation by an atom. First observed by Henri Becquerel ( ). Marie and Pierre Curie also studied it. Nobel Prize in 1903 (physics).

17 Studies of Natural Radioactivity Structure of the atom after Becquerel (early 1900s) Positively charged mass Electrons (-) Some atoms naturally emit one or more of the following types of radiation: alpha (α) radiation (later found to be He 2+ - helium nucleus) beta (β) radiation (later found to be electrons) gamma (γ) radiation (high energy light) γ α α γ Alpha particles Somehow gamma radiation is in there, too.

18 Radioactivity Three types of radiation were discovered by Ernest Rutherford: –  particles (positive, charge 2+, mass 7400 times of e - ) –  particles (negative, charge 1-) –  rays (high energy light)

19 Ernest Rutherford (1910) Scattering experiment: firing alpha particles at a gold foil

20 The Nuclear Atom Some alpha particles bounce off the gold foil. This means the mass of the atom must be concentrated in the center and is positively charged! Thompson’s model could not be correct.

21 Ernest Rutherford The Nucleus and the Proton Structure of the atom after Rutherford (1910) The mass is not spread evenly throughout the atom, but is concentrated in the center, the nucleus. Electrons (-) are now outside the nucleus. The positively charged particles in the nucleus are protons.

22 James Chadwick – The Neutron Structure of the atom after Chadwick (1932) In the nucleus with the protons are particles of similar mass but no electrical charge called neutrons. Electrons (-) are now outside the nucleus in quantized energy states called orbitals. (From Niels Bohr and quantum mechanics) The positively charged particles in the nucleus are protons. n n +

23 Structure of the Atom proton (+) neutron nucleus - responsible for the mass of the atom, positively charged electrons - responsible for the volume and size of the atom, negatively charged m m

24 Subatomic Particles Protons and electrons are the only particles that have a charge. Protons and neutrons have essentially the same mass. The mass of an electron is so small we ignore it.

25 Atomic Facts FeatureSizeMass 1 amu = 1 atomic mass unit = x g Electrons are outside the nucleus in quantized energy states called orbitals. proton (+) m amu neutron (0) m amu electron (-) m ??? amu + n n

26 Symbols of Elements Elements are symbolized by one or two letters.

27 Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z)

28 Atomic Mass The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

29 Atomic Number The number of protons in the nucleus is called the atomic number Z. Z determines the identity of an element. Saying “the atomic number of an element is 6” is the same as saying “carbon.” The number of electrons in the atom is also Z (because atoms have no net electric charge). How many neutrons are in C? Carbon atom - proton - neutron

30 12 C - proton - neutron Isotopes The number of protons and neutrons (nucleons) in an element is called the mass number A. How many neutrons are in C? The answer is “it depends on the isotope.” An element may have different numbers of neutrons but NOT different numbers of protons. Atoms of an element with different numbers of neutrons are called isotopes of that element. A = Z + number of neutrons. A 6 Z

31 Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons C 12 6 C 13 6 C 14 6 C

32 Isotopes number of protons (Z) number of neutrons mass number (A) number of electrons symbol C or C C or C O or O U or U

33 Atomic Masses Atomic masses are based on 12 C. The mass of 12 C (or C-12) is defined to be exactly 12 amu.

34 Atomic Masses The mass (weight) shown in the periodic table is the mass of the element as its occurs naturally. If the element has more than one isotope, the mass shown is the weighted average of the masses of the isotopes. Mg has 3 isotopes. 24 Mg 78.99% amu 25 Mg 10.00% amu 26 Mg 11.01% amu weighted average of Mg: x x x amu atomic weight of Mg based on natural abundance: amu

35 Ions Atoms can gain or lose electrons to become charged particles called ions. –A chemical particle that contains a positive or negative charge Cations are positively charged ions. –Formed when an atom loses electrons Anions are negatively charged ions. –Formed when an atom gains electrons

36 Ions Net charge = 0Net charge = +1 Formation of a cation 1p e-e- e-e- + Hydrogen atom 1p, 0 n, 1 e - Hydrogen ion (cation) 1p, 0 n, 0 e - 1 H 1 H +

37 Ions 8p 8n 8e - 8p 8n Oxygen atom 8p, 8 n, 8e - Oxygen ion (anion) 8p, 8n, 10e - 16 O 16 O 2- 10e - 2e - + Net charge = -2Net charge = 0 Formation of an anion

38 Mass NumberCharge Atomic Number Nuclear Symbols X Charge = # p - # e -

39 Nuclear Symbols Using nuclear symbols to determine the number of p, n, e, and total charge O 16 8 Mass Number = Atomic Number = 16 8 # protons = atomic number = 8 # neutrons = Mass # - Atomic # = = 8 # electrons = # protons = 8

40 Nuclear Symbols Mass Number = Atomic Number = # protons = atomic number = 8 # neutrons = Mass # - Atomic # = = 8 # electrons = # protons - charge = 8 - (-2) = O

41 Nuclear Symbols Ba Mass Number = Atomic Number = # protons = atomic number = 56 # neutrons = Mass # - Atomic # = = 81 # electrons = # protons - charge = 56 - (+2) = 54

Sn 50 Nuclear Symbols - Atoms Example: Write the nuclear symbol for the following atoms: 1)50 p, 70 n 2)17 e -, 20 n 37 Cl 17

43 Nuclear Symbols - Ions Practice writing nuclear symbols from information given: 1)53 p, 74 n, 54 e - 53 proton (= atomic number)  I 74 neutrons + 53 proton  mass number = electrons (one more than protons)  I 1- 53

44 2)23 e -, 30 n, net charge = +3 # protons? 23 electrons, but charge of 3+ ie 3 more protons than electrons  p= 26  Atomic number = 26  element = Fe Fe