Unit 2 - Lecture 1: Structure of the Atom Atomic Theory and what we understand today about… Atoms, Molecules, and Ions

The Power of 10 Have you ever wondered your placement in the universe and the things that make it up? Picture this! http://micro.magnet.fsu.edu/primer/java/scienceopticsu/powersof10/

A HISTORY OF THE STRUCTURE OF THE ATOM Unit 2 - Lecture 1: Structure of the Atom A HISTORY OF THE STRUCTURE OF THE ATOM

Let’s begin with… Greek Philosopher Democritus (460-370 B.C.): all matter composed of small atoms atomos = indivisible What did Democritus conclude about cutting matter in half? There was a limit to how far you could divide matter. You would eventually end up with a piece of matter that could not be cut. Why weren’t Democritus’s ideas accepted? Aristotle was a very famous Greek philosopher who believed that matter could be divided into smaller and smaller pieces forever. He held a very strong influence on popular belief and his views on this were accepted for two thousand years.

1803, John Dalton: atoms are the fundamental building blocks of matter He performed many experiments to study how elements join together to form new substances He found that they combine in specific ratios and he supposed it was because the elements are made of atoms.

Dalton's Postulates Each element is composed of extremely small particles called atoms. Atoms are indivisible and indestructible particles. Figure 2.1 John Dalton (1766-1844)

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. Figure 2.1 John Dalton (1766-1844)

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. Figure 2.1 John Dalton (1766-1844)

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.

John Dalton’s Atomic Theory (ca 1803) Unit 2 - Lecture 1: Structure of the Atom John Dalton’s Atomic Theory (ca 1803) Each element is composed of extremely small particles called atoms. All atoms of a given element are identical. The atoms of different elements are different and have different properties (including different masses). 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. 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 Definite Composition. 11

John Dalton’s Atomic Theory Unit 2 - Lecture 1: Structure of the Atom John Dalton’s Atomic Theory Almost right. A good start. very small Structure of the atom after Dalton (ca. 1810)

Unit 2 - Lecture 1: Structure of the Atom J.J. Thomson (1897): Cathode Rays Atoms subjected to high voltages give off cathode rays. I can demonstrate this for you!

JJ Thomson: https://youtu.be/oddjdB0qfMg What did he discover? A particle that has MASS and has a NEGATIVE CHARGE The mass was 2000x smaller than the smallest known atom (Hydrogen) This particle must be FROM an atom! Named it the electron

Unit 2 - Lecture 1: Structure of the Atom J.J. Thomson: Summarizing what he saw Thomson concluded that the negative charges came from within the atom. A particle smaller than an atom had to exist. The atom was divisible! Since the gas was known to be neutral, having no charge, he reasoned that there must be positively charged particles in the atom. But he could never find them. Electrons are in atoms.

Unit 2 - Lecture 1: Structure of the Atom J.J. Thomson – The Electron “Plum pudding” model: Negative electrons are embedded in a positively charged mass. Electrons (-) Unlike electrical charges attract, and that is what holds the atom together. Positively charged mass Structure of the atom after Thomson (ca. 1900)

Radioactivity Radioactivity is the spontaneous emission of radiation by an atom. First observed by Henri Becquerel (1852-1908). Marie and Pierre Curie also studied it. Nobel Prize in 1903 (physics).

Unit 2 - Lecture 1: Structure of the Atom Studies of Natural Radioactivity Some atoms naturally emit one or more of the following types of radiation: alpha (α) radiation (later found to be He2+ - helium nucleus) beta (β) radiation (later found to be electrons) gamma (γ) radiation (high energy light) α Alpha particles Electrons (-) γ γ Positively charged mass α Somehow gamma radiation is in there, too. Structure of the atom after Becquerel (early 1900s)

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) Figure 2.8

Unit 2 - Lecture 1: Structure of the Atom Ernest Rutherford (1910) Scattering experiment: Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil (2000 atoms thick)

The Nuclear Atom Most of the positively charged “bullets” passed right through the gold atoms in the sheet of gold foil without changing course at all. Some of the positively charged “bullets,” however, did bounce away from the gold sheet as if they had hit something solid. He knew that positive charges repel positive charges. Figure 2.11

Rutherford’s ideas This could only mean that the gold atoms in the sheet were mostly open space. Atoms were not a pudding filled with a positively charged material. Rutherford concluded that an atom had a small, dense, positively charged center that repelled his positively charged “bullets.” He called the center of the atom the “nucleus” The nucleus is tiny compared to the atom as a whole.

Unit 2 - Lecture 1: Structure of the Atom Ernest Rutherford The Nucleus and the Proton The mass is not spread evenly throughout the atom, but is concentrated in the center, the nucleus. The positively charged particles in the nucleus are protons. Electrons (-) are now outside the nucleus. Structure of the atom after Rutherford (1910)

Unit 2 - Lecture 1: Structure of the Atom James Chadwick – The Neutron 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 Structure of the atom after Chadwick (1932)

Unit 2 - Lecture 1: Structure of the Atom proton (+) neutron electrons - responsible for the volume and size of the atom, negatively charged 10-10 m nucleus - responsible for the mass of the atom, positively charged 10-14 m

If this stadium were the size of an atom’s electron cloud, the nucleus would be the size of a marble setting on the 50 yard line. Electrons occupy the VOLUME, protons and neutrons constitute the MASS of an atom.

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 (2000 time smaller than the proton), we ignore it. Table 2.1 0.0005486 is the way to write this number!

Unit 2 - Lecture 1: Structure of the Atom Atomic Facts Feature Size Mass 1 amu = 1 atomic mass unit = 1.66054 x 10-24 g electron (-) 10-18 m ??? 0.0006 amu + n n proton (+) 10-15 m 1.0073 amu Electrons are outside the nucleus in quantized energy states called orbitals. neutron (0) 10-15 m 1.0087 amu

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

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

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

Unit 2 - Lecture 1: Structure of the Atom Atomic Number Carbon atom 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? - proton - neutron

Unit 2 - Lecture 1: Structure of the Atom Isotopes A 12C Z 6 The number of protons and neutrons in an element is called the mass number A. A = Z + number of neutrons. 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. - proton - neutron How many neutrons are in C? The answer is “it depends on the isotope.”

Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C

Unit 2 - Lecture 1: Structure of the Atom Isotopes number of protons (Z) number of neutrons 6 8 92 146 mass number (A) number of electrons symbol 12 6 12C or C-12 6 14 6 14C or C-14 6 16 8 16O or O-16 8 238 92 238U or U-238 92

Unit 2 - Lecture 1: Structure of the Atom Atomic Masses Atomic masses are based on 12C. The mass of 12C (or C-12) is defined to be exactly 12 amu.

Unit 2 - Lecture 1: Structure of the Atom 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. 24Mg 78.99% 23.985 amu 25Mg 10.00% 24.986 amu 26Mg 11.01% 25.983 amu weighted average of Mg: 0.7899x23.985 18.946 0.1000x24.986 2.499 0.1101x25.983 +2.861 24.31 amu atomic weight of Mg based on natural abundance: 24.31 amu

Unit 2 - Lecture 1: Structure of the Atom 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

Unit 2 - Lecture 1: Structure of the Atom Ions Formation of a cation 1p e- + Hydrogen atom 1p, 0 n, 1 e- Hydrogen ion (cation) 1p, 0 n, 0 e- 1 H 1 H+ Net charge = 0 Net charge = +1

Unit 2 - Lecture 1: Structure of the Atom Ions Formation of an anion 8p 8n 8e- + 2e- 8p 8n 10e- Oxygen atom 8p, 8 n, 8e- Oxygen ion (anion) 8p, 8n, 10e- 16 O2- 16 O Net charge = 0 Net charge = -2

Unit 2 - Lecture 1: Structure of the Atom Nuclear Symbols Mass Number Charge Atomic Number X Charge = # p - # e-

Unit 2 - Lecture 1: Structure of the Atom 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 # = 16 - 8 = 8 # electrons = # protons = 8

Unit 2 - Lecture 1: Structure of the Atom Nuclear Symbols O 16 8 2- Mass Number = Atomic Number = 16 8 # protons = atomic number = 8 # neutrons = Mass # - Atomic # = 16 - 8 = 8 # electrons = # protons - charge = 8 - (-2) = 10

Unit 2 - Lecture 1: Structure of the Atom Nuclear Symbols Ba 137 56 2+ Mass Number = Atomic Number = 137 56 # protons = atomic number = 56 # neutrons = Mass # - Atomic # = 137 - 56 = 81 # electrons = # protons - charge = 56 - (+2) = 54

Nuclear Symbols - Atoms Unit 2 - Lecture 1: Structure of the Atom Nuclear Symbols - Atoms Example: Write the nuclear symbol for the following atoms: 1) 50 p, 70 n 2) 17 e-, 20 n 120Sn 50 37Cl 17

Unit 2 - Lecture 1: Structure of the Atom Nuclear Symbols - Ions Practice writing nuclear symbols from information given: 53 p, 74 n, 54 e- 53 proton (= atomic number)  I 74 neutrons + 53 proton  mass number = 127 54 electrons (one more than protons)  1- 127I1- 53

2) 23 e-, 30 n, net charge = +3 # protons 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 56 3+ Fe 26

Our classmates are going to present information about electrons/energy levels and more and that is how we will learn more on this timeline of the atomic theory.