3 I. Law vs. Theory 1) Scientific law = a generalization of scientific observations that describes what happens (does not explain) 2) Theory (model) = a set of assumptions used to explain observations and predict new observations
4 THEORY: a) Can never be truly proven 100% correct. b) Inevitably change (and must sometimes be abandoned) as more information becomes available. c) Considered successful when they explain observations and, more importantly, predict new observations.
6 Atomic Structure 1) EARLY ATOMIC THEORY: a) 400 B.C- Democritus (Greek philosopher) Coined the term atom (meaning indivisible) to describe the smallest particles of matter Did not experiment
7 2) JOHN DALTON (EARLY 1800 S): a) Father of the Modern Atomic Theory b) Dalton s Atomic Theory helped explain measurable observations and successfully predict new observations (while stimulating additional research from Dalton and other scientists).
8 John Dalton c) Assumptions of Dalton s Atomic Theory (proposed by Dalton in 1803): I. All matter is composed of atoms II. All atoms of the same element are identical in size, mass, other properties. Atoms of different elements differ in size, mass, and other properties.
9 John Dalton III. Atoms cannot be subdivided, created, or destroyed IV. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. V. In chemical reactions, atoms are combined, separated, or rearranged.
10 John Dalton d) Dalton used his Atomic Theory to help correctly predict new observations leading to the Law of Multiple Proportions.
11 John Dalton e) Law of Multiple Proportions = The masses of one element that combine with a constant mass of another element to form more than one compound are in the ratio of small whole numbers Carbon monoxide and carbon dioxide contain oxygen in a 1:2 ratio. 12 g of carbon reacts with 16 g of oxygen to form carbon monoxide, CO. 12 g of carbon will react with 32 g of oxygen to form carbon dioxide, CO 2.
12 John Dalton f) Dalton s Atomic Theory explained many observations and correctly predicted many additional observations. It DID NOT correctly predict all new observations (no theory ever does).
13 g) Dalton s Atomic Theory has been modified to explain the new observations: Atoms of one element can have different masses (isotopes). Atoms can be subdivided (but not in a chemical reaction or physical change) Nuclear reactions An atom of one element can be changed into an atom of another element (but not in a chemical reaction or physical change) Nuclear reactions
14 John Dalton: g) As with all theories, Atomic Theory has been changed and expanded over time (and will continue to change and expand) to explain new observations.
15 3) CATHODE-RAY TUBE EXPERIMENTS(late 1800 s): Cathode-ray tube = A glass tube containing a gas at a very low pressure which contains a negative electrode (cathode) and a positive electrode (anode). When high voltage is applied, a cathode ray passes from the cathode to the anode causing the low pressure gas in the tube to glow (different gases glow different colors).
16 3) CATHODE-RAY TUBE EXPERIMENTS (late 1800 s): a) Early cathode-ray tube experiments: Proposed that a cathode ray consists of tiny particles with mass and they are negatively charged.
17 3) CATHODE-RAY TUBE EXPERIMENTS (late 1800 s): b) J. J. Thomson discovered the electron Used early cathode-ray tube experiments and some of his own findings to support the hypothesis that electrons are negatively charged particles
18 J.J. THOMSON: Used a cathode ray tube- as voltage across the tube was increased, a beam of light (cathode ray) became visible Cathode ray = beam of electrons seen b/c of excited gas
19 Found that the beam was deflected by both magnetic and electrical fields OBSERVATION: Noticed that the cathode rays were attracted to the positive electrode, called the anode. J.J. THOMSON:
20 What conclusions did Thompson have? The cathode rays were made up a very small negatively charged particle ELECTRONS
21 J.J. THOMSON: OBSERVATION: Measured the bending of the path of the cathode rays and was able to determining the ratio of an electron s charge to its mass. Found that the ratio was always the same, regardless of the metal used
22 What conclusions did Thompson have? Concluded that the negatively charged particles were much lighter than the lightest know atom (hydrogen), which meant that atoms had a structure!
23 c) Later Protons were also discovered using a modified cathode-ray tube
24 d) Thomson s Model of the Atom (Plum Pudding Model) Thomson postulated that an atom consisted of a diffuse cloud of positive charge with the negative electrons embedded randomly in it. This represented a major change from Dalton s model of atoms as indivisible.
25 The existence of the electron raised new questions: if electrons are part of all matter and they possess a negative charge, how can all matter be neutral? Also, if the mass of an electron is so small, what accounts for the rest of the mass in a typical atom? More experiments to come
26 4) ERNEST RUTHERFORD & THE GOLD FOIL EXPERIMENT a) Rutherford discovered the NUCLEUS using alpha particles & gold foil Alpha particle = a relatively large positively charged particle
27 ERNEST RUTHERFORD b) Gold-Foil Experiment (around ): Thin gold foil was bombarded by alpha particles and the path of the alpha particles was charted after they passed through the gold foil.
28 ERNEST RUTHERFORD c)expected results: The massive alpha particles (positively charged) were expected to crash through the gold foil with little or no deflections; there was nothing in Thomson s model of the atom to cause anything more than minor deflections of the alpha particles.
29 ERNEST RUTHERFORD d) OBSERVATION: Some alpha particles were deflected at large angles, and some were redirected backward (surprising results!) What conclusion did Rutherford draw from this evidence? The POSITIVE particles of the atom mustnot be spread out evenly, but instead must be concentrated at the center of the atom-the(nucleus).
30 ERNEST RUTHERFORD e) OBSERVATION: Most of the alpha particles passed through the gold foil with few deflections. What conclusion did Rutherford draw from this evidence? Most of the alpha particles did not hit anything and passed straight through the gold atoms so therefore, most of the volume of an atom consists of empty space.
31 ERNEST RUTHERFORD f) Rutherford s Nuclear Model Explained the neutral nature of matter: the positive charge of the nucleus balances the negative charge of the electrons.
32 ERNST RUTHERFORD Suggested that the electrons travel around the positively charged nucleus. The early nuclear model did not account for all the atom s mass By 1920, Rutherford refined the concept of the nucleus and concluded that the nucleus contained positively charged particles called protons.
33 ERNEST RUTHERFORD r14.swf Animation on NEXT slide!
35 (a) When a beam of alpha particles is directed at a thin gold foil, most particles pass through the foil undeflected, but a small number are deflected at large angles and a few bounce back toward the particle source. (b) A closeup view shows how most of an atom is empty space and only the alpha particles that strike a nucleus are deflected.
36 5) JAMES CHADWICK (1932) a) Discovered the neutron
38 1) WHAT IS THE DIFFERENCE BETWEEN AN ATOM AND AN ELEMENT? a) Element = a substance that cannot be broken down to other substances by a chemical reaction. b) Atom: the smallest particle of an element that can exist either alone or in combination with other atoms.
39 2) 3 SUBATOMIC PARTICLES OF AN ATOM
40 2) SUBATOMIC PARTICLES OF AN ATOM a) Subatomic particle = a particle smaller than an atom Ex: proton, neutron, electron
41 2) SUBATOMIC PARTICLES OF AN ATOM b) An atom is composed of subatomic particles including protons, neutrons, and electrons (plus scientists have determined that protons and neutrons have their own structures and they are composed of quarks- these particles will not be covered since scientists do not yet understand if or how they affect chemical behavior).
42 c) Summary of subatomic particles Particle Symbol Location Electron e- Proton Neutron p n In the space surrounding the nucleus In the nucleus In the nucleus Relative Electron Charge Relative Mass Actual Mass (g) -1 1/ g g g
43 2) SUBATOMIC PARTICLES OF AN ATOM d) Protons give the nucleus the positive charge & determines the identity of an atom e) Protons and neutrons have about the same mass and are over 1840 times more massive than electrons.
44 2) SUBATOMIC PARTICLES OF AN ATOM f) Most of the mass of an atom is located in the nucleus (protons & neutrons). (Nuclear forces hold the particles of a nucleus together. These forces only act over a very short range.) g) Atoms are electrically neutral b/c of the number of protons EQUALS the number of electrons.
45 2) SUBATOMIC PARTICLES OF AN ATOM h) Nucleus = the positively charged, dense central portion of an atoms Contains nearly all of the atom s mass, but takes up a very small fraction of its volume
46 3) ATOMIC NUMBER & AVERAGE ATOMIC MASS 19 K Potassium Atomic Number Average Atomic Mass Atomic number: the number of protons in the nuclei Also represents the number of electrons that an atom has since atoms are electrically neutral Ex: Atomic number of Na is 11; Na has 11 protons and 11 electrons
47 3) ATOMIC NUMBER & AVERAGE ATOMIC MASS Average Atomic Mass: the weighted average of the atomic masses of naturally occurring isotopes of an element. Isotopes = atoms of the same element with different number of NEUTRONS, therefore, different masses. Periodic Table lists average atomic mass
48 Weighted averages account for the percentages of each isotope of a given element Important b/c they indicate relative mass relationships in chemical reactions
49 3) ATOMIC NUMBER & AVERAGE ATOMIC MASS Ex: How average atomic mass is calculated
50 4) ISOTOPES a) Isotopes = atoms of the same element with different number of neutrons, therefore, different masses. Isotopes of an element all have the SAME NUMBER of protons and electrons Named by their mass numbers
51 4) ISOTOPES Mass number = the total number of protons and neutrons in the nucleus of an isotope. Mass Number = # Protons + # Neutrons
52 4) ISOTOPES Ex: H-1 H-2 H-3
53 4) ISOTOPES Example 1: ALL carbon atoms have how many protons? 6 (atomic number) Most carbon atoms have 6 neutrons. What is their mass number? 12
54 10) ISOTOPES Some carbon atoms have 8 neutrons. What is their mass number? 14 C-12 and C-14 are isotopes of carbon Both isotopes have 6 electrons & 6 protons, but differ in the # of NEUTRONS!!
55 10) ISOTOPES Example 2: How many neutrons are in a sodium-23 atom? 12 Example 3: How many protons, electrons and neutrons are there in an atom of chlorine-37? 17 protons; 17 electrons (atomic number is 17) 37 (protons + neutrons) 17 protons = 20 neutrons
56 Hyphen Notation: 10) ISOTOPES Write the hyphen notation for a hydrogen isotope with a mass number of 3. H-3 or Hydrogen-3 Nuclear Symbol Notation Write the nuclear symbol for a hydrogen isotope with a mass number of 3.
57 Complete the following table. Remember: atomic number = number of protons = number of electrons mass number = atomic number + number of neutrons
58 Nuclear Symbol Atomic Number Mass Number Number of Protons Number of Neutrons Number of Electrons
59 THE MOLE = SI unit for amount of substance
60 The Mole: A mole simply represents a counting unit, much in the same way that a dozen represents a set of twelve. Ex: Just as you can have a dozen cans of soda or a dozen donuts, you can have a mole of stars or a mole of water molecules. So a dozen eggs represents 12 eggs, a mole of carbon represents X carbon atoms X is also called Avogadro s Number = the number of particles in exactly one mole of a pure substance.
61 The Mole: There are 6.02 X atomic mass units (amu) in one gram!! This conversion allows us to use the masses listed on the periodic table for both atomic mass and molar mass. Atomic mass = is the mass of one atom, is measured in atomic mass units (amu) Mass of C-12 is arbitrary assigned a mass of 12 amu 1/12 the mass of a carbon-12 atom is 1 amu Mass of any atom is expressed relative to the mass of one atom of carbon-12
62 The Mole: Molar mass = is the mass of one mole of atoms or molecules, is measured in grams. Numerically equal to the atomic mass 1 mole of a substance is 1 molar mass of that substance Ex: 1 mole of nitrogen is equal to g of nitrogen!
63 Atomic & Molar Masses Element & Symbol Atomic Mass Mass of 1 Atom Molar Mass Mass of 6.02 X atoms Carbon (C) amu g Helium (He) amu g Copper (Cu) amu g Potassium (K) amu g
64 The Mole: REMEMBER:one mole of atoms contains 6.02 X atoms. One mole of molecules contains 6.02 X molecules. One mole of formula units contains 6.02 X formula units. One mole of ions contains 6.02 X ions.
65 Conversions: Using Charts! Ex: Conversion factors: 12 inches = 1 foot 3 feet = 1 yard Set up conversion charts so your units will cancel out! You multiple across the top of the chart and divide by what is on the bottom of the chart.
66 How many feet are in yards?
67 How many yards are in feet?
68 How many inches are in feet?
69 How many feet are in 1,743.2 inches?
70 NOTE: Mrs. H s ROUNDING We will use 6.02 X atoms We will round the molar mass values from the Periodic Table (0.5 + round to the whole #) Ex: Phosphours30.974g would be rounded to 31g
71 MOLESMASS MASS (grams) KEY IDEA: It is important to know the following conversion: The molar mass of an element is equal to the mass of 1 moleof atoms of that element. Example: 1 mole of zinc is equal to grams of zinc! Use the factor label method (units will cancel out!)
72 MOLESMASS MASS (grams) Ex: What is the mass in grams of 3.5 mol of the element copper (Cu)? 3.5 mol Cu 64 g Cu 1 mol Cu = 224 g Cu
73 1. What is the mass in grams of 2.25 mol of iron (Fe)?
74 2. What is the mass in grams of mol of K?
75 MASS (grams) MOLES 1 mol of an element = the molar mass of that element
76 MASS (grams) MOLES Ex: A chemist produced 11.9g of Al. How many moles of Al has been produced? 11.9 g Al 1 mol Al 27 g Al = mol Al
77 1. How many moles of Ca are contained in 5.0g of Ca?
78 2. How many moles of Au are contained in 3.6 x g of Au?
79 ATOMS MOLES 1 mol of an element = 6.02 x atomsof that element
80 ATOMS MOLES Ex: How many moles of Ag are in 3.01 x atoms? 3.01 x Ag atoms 1 mol Ag = 0.5 mol Ag 6.02 x Ag atoms
81 1. How many moles of Pb are equivalent to 1.5 x atoms?
82 2. How many moles of tin are equivalent to 2500 atoms?
83 3. How many atoms of Al are contained in 2.75 mol?
85 The Mole, Molar Mass, # of Atoms Practice Problems What is the conversion factor between moles and atoms? 1 mole = 6.02 X10 23 atoms
86 The Mole, Molar Mass, # of Atoms Practice Problems What is the conversion factor between moles and mass? 1 mole = molar mass (grams)