Section 3.1 Substances Are Made of Atoms

Transcription

1 Section 3.1 Substances Are Made of Atoms Objectives: 1. State the three laws that support the existence of atoms. 2. List the five principles of John Dalton s atomic theory. Vocabulary: law of definite proportions law of conservation of mass law of multiple proportions Atomic Theory Atomic theory, which states that atoms are the building blocks of matter, has been around as long ago as 400 BCE. However, scientific experiments supporting the existence of atoms did not appear until the late 17 th century. Experimental results show that the following laws support the current atomic theory: Law of Definite Proportions each specific chemical compound always contains exactly the same elements in exactly the same proportions by number, or by mass (or by weight), or by volume. Law of Conservation of Mass mass cannot be created or destroyed in normal chemical reactions (the mass of the reactants equals the mass of the products). Law of Multiple Proportions atoms combine in whole number ratios. Dalton s Atomic Theory In 1808, John Dalton, an English high school teacher, used the Greek concept of the atom and the law of definite proportions, the law of conservation of mass, and the law of multiple proportions to develop an atomic theory. According to Dalton, elements are composed of one type of atom and compounds are composed of two or more types of atoms. Dalton s atomic theory can be summarized by the following statements: 1. All matter consists of atoms that cannot be divided, created, or destroyed. 2. Atoms of the same element are identical in their physical and chemical properties. 3. Atoms of different elements are chemically and physically different. 4. Atoms of different elements combine in simple, whole number ratios to form compounds. 5. In chemical reactions, the atoms in compounds will separate, rearrange, and combine but atoms are never created, destroyed, or changed. The huge advance in human thinking about these statements that these propositions can be tested. Note that propositions 1 and 2 are no longer true. YouTube (2 min) (4 min) 1

2 Section 3.2 Structure of Atoms Objectives: 1. Describe the evidence for the existence of electrons, protons, and neutrons, and describe the properties of these subatomic particles. 2. Discuss atoms of different elements in terms of their numbers of electrons, protons, and neutrons, and define the terms atomic number and mass number. 3. Define isotope, and determine the number of particles in the nucleus of an isotope. Vocabulary: electron nucleus proton neutron atomic number mass number isotope charge Historical Development of Atomic Theory Experiments were conducted in the mid-19 th to early 20 th century that led to the discovery that atoms are themselves made of electrons, protons, and neutrons. J.J. Thomson discovered the electron using a cathode ray tube Electrons have negative charge because they are emitted from the negatively charged cathode. Thomson also placed a paddle wheel in the tube and it turned, which indicated that electrons are particles with mass. Key hypothesis: atoms have neutral charge, so there must exist positively charged subatomic stuff that balanced the negativity of the electrons. Thomson proposed his Plum Pudding Model of the atom. 2

3 Ernest Rutherford s Gold Foil Experiment According to the plum pudding model, all alpha particles should go straight through. But some were deflected at backward angles! It was like firing a 15 inch shell at paper tissue and it bounced back to hit you. Using the angle of deflection, scientists determined: A heavy, positively charged nucleus much larger than the alpha particle. The radius of the atom is 10,000 times larger than that of the alpha particle. The atom is mostly empty space! It consists of a tiny, very dense nucleus surrounded by distant electrons. The nucleus contains protons (p + ) that are 2000 times heavier than the electron (e - ). The nucleus also contains neutrons (n o ) that are neutrally charged and mass the same as a proton. (Irene Joliot-Curie, daughter of Marie Curie, was a co-discoverer of neutrons when alpha particles hit barium). Even though protons have positive charge, they form a stable nucleus! Neutrons hold protons together in the nucleus. All atoms with 2 or more protons in the nucleus also have neutrons in the nucleus. Atomic Number and Mass Number Name Symbol Charge Common Charge Notation Mass Common Mass Notation Electron e x10-19 coulombs x10-31 kg 0 Proton p x10-19 coulombs x10-27 kg 1 Neutron n o 0 coulombs x10-27 kg 1 Mass Number = total number of protons and neutrons in the nucleus. Knowing a mass number does not help identify an element since the number of neutrons can change for the same element. (Also called the atomic mass.) 26 Fe Atomic Number = number of protons in the nucleus. The atomic number is unique to each element; it defines the type of element. 3

4 Isotopes The nucleus of the same element can have a different number of neutrons. Atoms of the same element (equal number of protons in the nucleus) that have a different number of neutrons are called isotopes. For example, there are two isotopes of helium: helium-3 and helium-4. All isotopes of an element have the same atomic number Atomic mass numbers of isotopes of the same element are not the same because the numbers of neutrons are different. The mass number listed in the periodic table is an average of the mass numbers of the naturally occurring isotopes of the element. Name Symbol # neutrons # protons Mass # Abundance Lead Pb % Lead Pb % Lead Pb % Lead Pb % The mass number used in the periodic table is determined as follows: Mass = Σ (Abundance of Isotope)(Mass # of Isotope) For example, the average atomic mass of the naturally occurring isotopes of lead is: Mass # = (1.4%)(204) + (24.1%)(206) + (22.1%)(207) + (52.4%)(208) = How do electrons stay close to the nucleus? Electromagnetic force causes the negatively charged electrons to be attracted to the positively charged nucleus. The equation for this attraction is called Coulomb s Law: Electromagnetic Force (F) = charge on one particle (q 1) x charge on other particle (q 2 ) (distance between charged particles) 2 = q 1x q 2 d 2 How can proton form a stable nucleus? Protons have a positive charge and exist together in the nucleus. The nature of magnetic force predicts that they should repel! However, at very small distances the strong nuclear force overcomes Coulombic repulsion and holds protons together. In addition, neutrons stabilize the nucleus neutrons are like glue holding the nucleus together. YouTube History of Atoms (Bozeman, 9:10) 4

5 Atom are mostly empty space-bill Nye (6:37) Development of the Atomic Theory (9:52) Cartoon of the History of the Atom Theory Chemistry Atomic Theory Timeline (2:07) Have you ever seen an atom? (Nature Video, 2:32) Early Atomic Theory: Dalton, Thomson, Rutherford, and Millikan Section 3.3: Electron Configuration Objectives: 1. Compare the Rutherford, Bohr, and quantum models of the atom. 2. Explain how the wavelengths of light emitted by an atom provide evidence for the modern model of the atom and information about electron energy levels. 3. List the four quantum numbers, and describe their significance. 4. Write the electron configuration of an atom by using the Pauli Exclusion Principle and the Aufbau Principle. Vocabulary: orbital electromagnetic spectrum ground state excited state quantum number Pauli exclusion principle electron configuration aufbau principle Hund s rule Atomic Models Building a model helps scientists imagine what may be happening at the microscopic level. Models have limitations and have to be modified or discarded as new information comes available. Dolton s Model

6 Thomson s Model Key hypothesis: atoms have neutral charge, so there must exist positively charged subatomic particles that balance the negativity of the electrons. Thomson proposed his Plum Pudding Model of the atom. Rutherford s Model Rutherford s gold foil experiment led replacement of the Plum Pudding model with the nuclear model of the atom. Rutherford suggested that electrons revolved around the nucleus like planets orbited around the sun. But this did not explain why the negative electrons did not get pulled into the positive nucleus, because opposite charges attract. 6

7 Bohr s Model Electrons can only exist in energy levels (also called orbitals) that are specific distances from the nucleus. The electrons do not exist in between the energy levels. Energy levels closer the nucleus are lower in energy. The energy levels further from the nucleus are higher in energy. Normally an electron hangs out in the lowest possible energy level, called the ground state level. If an atom absorbs energy, the electron can jump to a higher energy level. As the atom cools down, the excited electron falls into a lower orbital, loses energy, causing the release of light energy as a photon, or light packet. The color of the released light is determined by how far the electron falls and how much energy it gives up. Energy Transition Calculations Using Bohr s Model d c e f g b a h i relative energies The location of an electron in an energy level is called its energy state. This diagram shows a cross-section of the transitions that an electron can make from a higher (excited state) to a lower (lower excited state or ground state) energy level. The lower energy levels are closer to the nucleus. The type of light given off when an electron drops from a higher to a lower level is determined by the difference in the relative energies of the energy levels. For example, the energy released when an electron falls from the level having a relative energy of 31 to the ground state level (indicated by arrow d) is 31 1 = 30. This energy transition might correspond to the color red. Quantum Mechanics: The Structure of Atom (6:11 min) 7

8 Electromagnetic Spectrum Line Spectrum When a particular atom absorbs energy, its electrons are elevated from their ground state into higher or excited states. As the atom cools, the excited electrons fall into lower energy states, releasing light in the process. Each element has a unique set of energy levels, and gives off a unique set of colors called a line spectrum. The line spectrum can be used to identify the type of element, similar to using a fingerprint to identify a person. The following diagram shows the emission line spectrum for hydrogen, mercury, and neon. Also shown is the absorption line spectrum for hydrogen. 8

9 Modern Model Quantum View of Atomic Structure The quantum model is the present-day view of the atom. The quantum model places electrons in one of several types of orbitals, or regions of high probability for finding a particular electron. It is not possible to know exactly where the electron is or the direction it is going. The following diagram attempts to show a three dimensional view of the four types of orbitals, the s, p, d and f. Note that the s orbital has one sublevel the p orbital has three sublevels, the d orbital has five sublevels, and the f orbital has seven sublevels. Each sublevel can hold a maximum of two electrons. See: Electron Placement Electrons are placed into a particular orbital based on a set of quantum numbers that are unique to each electron in the atom. The principal quantum number (n) indicates the main energy level occupied by the electron and divides the periodic table into rows. n can have values of 1, 2, 3, 4, 5, 6, and 7. Higher n values are further from the nucleus and have higher levels of energy. The principal quantum number is further divided into the l, m, and spin sublevels. Angular momentum quantum number (l) indicates the shape of the orbital and can have values of 1, 2, 3, Magnetic quantum number (m) is a subset of l and indicates the orientation and can have values of, -2, -1, 0, 1, 2, Spin quantum number indicates the spin (+½, -½ or and ) of the electrons magnetic field and can hold a maximum of two electrons 9

10 Electron Placement Rules Aufbau Principal: electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states Pauli Exclusion Principal: An orbital can hold 0, 1, or 2 electrons only, and if there are two electrons in the orbital, they must have opposite (paired) spins. incorrect; electrons must spin in opposite directions correct; the unpaired electrons have the same spin Hund s Rule: If multiple orbitals of the same energy are available, unoccupied orbitals will be filled before occupied orbitals are reused (by electrons having different spins). Electron Configuration The arrangement of electrons in a particular atom can be found by determining the electron configuration of the electrons. The electron configuration can be determined filling the orbitals according to the yellow brick road in the following schematic: For example, the neutral sulfur atom has 16 electrons and thus has a configuration of: S: 1s 2 2s 2 2p 6 3s 2 3p 4 Each element s configuration builds on the previous elements configuration. To save space, use the configuration of the previous noble gas. Ne has 10 electrons, so: S: [Ne] 3s 2 3p 4 Taking this process to the extreme, an atom of the element with atomic number 118 has 118 electrons. By following the yellow brick road, the element would have the following electron diagram. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6 Orbital Diagrams The orbital diagram shows the energy levels and represents electrons by arrows placed inside the sub orbitals represented by boxes. The orbital diagram of sulfur is shown below. 10

11 The orbital diagram and the electron configuration process governed by the following rules: 1. No two electrons can be in the same place (Pauli Exclusion Principle) 2. Electrons assume the lowest possible energy level (aufbau principle) 3. Each individual orbital holds a maximum of two oppositely spinning electrons (Hund s Rule) 4. Given a choice, electrons prefer to occupy all available orbitals before pairing up. For 4 electrons in p orbitals, this means rather than Section 3.4: Counting Atoms Objectives: 1. Compare the quantities and units for atomic mass with those for molar mass. 2. Define mole and explain why this unit is used to count atoms. 3. Calculate either mass with molar mass or number with Avogadro s number given an amount in moles. Vocabulary: atomic mass mole molar mass Avogadro s number Atomic Mass Average mass of one copper atom: 11

12 1 atom Cu = x g Mass of one (pre-1982) penny is 3.13 g Cu. Therefore, the number of atoms in a copper penny is: (3.13 g Cu) ( 1 atom Cu x g Cu ) = 2.97 x 1022 Number of stars in the universe = = There are more atoms of Cu in a penny than stars in the universe! Using grams for the mass of atoms is inconvenient so chemists use the atomic mass unit (amu). One atomic mass unit = 1/12 th the mass of a carbon-12 atom (a carbon atom with 6 protons and 6 neutrons). One amu = x g. Introduction to the Mole Samples of elements have great numbers of atoms. To make working with these numbers easier, chemists use a unit called a mole (mol). A mole is defined as the number of atoms in exactly 12 grams of carbon-12. This number equals 6.02 x and is called Avogadro s number. Like the dozen, the mole is used to count things. The molar mass is the mass in grams of one mole (or 6.02 x atoms) of the element. Molar mass has the units of grams per mole (g/mol). The molar mass is equal to the atomic mass of an element. The molar mass of an element is used to convert between moles and grams.? mol (amount)x? g =? g (mass) mol? g (mass) x 1 mol? g =? mol (amount) Avogadro s number can be used to convert between moles and number of atoms.? mol (amount)x 6.02 x 1023 atoms 1 mol =? atoms (number of atoms)? atoms (number of atoms) x YouTube 1 mol 6.02 x atoms =? mol (amount) How big is a mole? (Not the animal, the other one.) (4:33 min) 12

13 1-1 The Mole & Avogadro s Number (4:19 min) Mr. Mole Man (0:43 min) A mole is a unit (2:27 min) Happy Mole Day to You (2:17 min) Happy Mole Day (Party in the USA) (2:34 min) Michael Offutt-A Mole is a Unit-Music Video (3:18 min) 13