Chapter 3 Atoms.

Chapter 3 Atoms

Vocabulary law of conservation of mass law of definite proportions
law of multiple proportions

Ch. 3 Section 1 The Atom: From Philosophical Idea to Scientific Theory
Objectives Explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. Summarize the five essential points of Dalton’s atomic theory. Explain the relationship between Dalton’s atomic theory and the law of conservation of mass, the law of definite proportions, and the law of multiple proportions.

Section 3.1 Defining the Atom
Around 460 B.C. Greek philosopher, Democritus, developed the idea of atoms. He asked this question: If you break a piece of matter in half, and then break it in half again, how many breaks will you have to make before you can break it no further? Democritus thought that it ended at some point, a smallest possible bit of matter. He called these basic matter particles, atoms.

Democritus’s Atomic Philosophy
Atoms were indivisible and indestructible Atom comes from the Greek word atomos, meaning “indivisible” His ideas: were based on philosophy didn’t explain chemical behavior lacked experimental support

Law of conservation of mass
A. Three Atomic Laws Rules that explain how matter interacts in chemical reactions: 1790’s- improved balances revolutionized quantitative analysis of chemical reactions.. allowed accurate measure of masses of elements and compounds Law of conservation of mass mass is neither created nor destroyed during ordinary chemical reactions or physical changes

A. Three Atomic Laws 2. Law of definite proportions
Rules that explain how matter interacts in chemical reactions: 2. Law of definite proportions a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source 3. Law of multiple proportions two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

Law of Conservation of Mass
Chapter 3 Section 1 The Atom: From Philosophical Idea to Scientific Theory Law of Conservation of Mass

Law of Multiple Proportions
Chapter 3 Section 1 The Atom: From Philosophical Idea to Scientific Theory Law of Multiple Proportions

John Dalton ( ) Dalton transformed Democritus’s ideas on atoms into a scientific theory. Dalton used experimental methods not just philosophy Studied ratio in which elements combine in chemical reactions Developed an atomic theory based on results of experiments Note: Dalton was always very smart and at the age of twelve he took over a Quaker school in Cumberland, England. Designated much of his life to researching color blindness because both he and his brother suffered from it.

Dalton’s Elements and Compounds

B. Dalton’s Atomic Theory (p. 64)
All elements are composed of extremely small particles called atoms. Atoms of an element are identical in size, mass and other properties; atoms of different elements differ in size, mass, and other properties.

B. Dalton’s Atomic Theory (p. 64)
3. Atoms cannot be subdivided, created, or destroyed 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged. (law of conservation of mass) (law of definite proportions and multiple proportions) (law of conservation of mass)

Scanning Tunneling Microscope (STM)
Gerd Binnig Dr. Heinrich Rohrer Invented in 1981 by Gerd Binnig and Heinrich Rohrer working for IBM in Zurich, Switzerland. Won 1986 Nobel Prize in physics for their discovery Opened door to field of Nanotechnology Scanning Tunneling Microscopy

Use of Scanning Tunneling Microscopy by IBM
Moving Atoms: Making the World's Smallest Movie The Boy and His Atom

Scanning Tunneling Microscope (STM)
STM provides ability to look at individual atoms detects the electrons of surface atoms maps the position and spacing of atoms

Dalton’s Atomic Theory has not been discarded.. only modified!
Modern Atomic Theory Not all aspects of Dalton’s atomic theory have proven to be correct. We now know that: Atoms are divisible into even smaller particles. A given element can have atoms with different masses (isotopes). Some important concepts remain unchanged: All matter is composed of atoms. Atoms of any one element differ in properties from atoms of another element. Dalton’s Atomic Theory has not been discarded.. only modified!

Chapter 3 Atoms

Vocabulary atom nuclear force

Chapter 3 Section 2 The Structure of the Atom Objectives Summarize the observed properties of cathode rays that led to the discovery of the electron. Summarize the experiment carried out by Rutherford and his co-workers that led to the discovery of the nucleus. List the properties of protons, neutrons, and electrons. Define atom. Describe nuclear forces

Section 3.2 Structure of the Atom
One change to Dalton’s atomic theory - atoms ARE divisible Into subatomic particles: Electrons, protons, and neutrons

Thomson’s Cathode-Ray Tube Experiment

Discovery of Electron- J.J. Thomson p. 69
J.J. Thomson used cathode ray tube to show presence of negatively charged particles...electrons Note: Called the negatively charged particles “corpuscles” JJ Thomson's Cathode-Ray Tube Experiment

Thomson’s Atomic Model
J. J. Thomson cathode ray experiment showed presence of electrons “plum pudding” model - believed e- like plums in + charged “pudding,” charge-to-mass ratio of an electron

Joseph John “J.J.” Thomson 1856-1940
Famous for discovery of the first subatomic particle, the electron, and for work on the atomic model. It was soon superseded by his student Ernest Rutherford’s nuclear model. 1906 received Nobel Prize for physics for his research into the electrical conductivity of gases. Thomson was a great teacher and an outstanding scientist. Seven of his students and assistants also received Nobel Prizes for their work.

CRT - Cathode Ray Tube Television Computer Monitor
Cathode ray tube (CRT) is a vacuum tube containing an electron gun (source of electrons) and fluorescent screen used to view images. It accelerates and deflects the electron beam onto fluorescent screen to create images. Can be found in everyday devises such as televisions, video games, computers, video cameras, monitors, automated teller machines, oscilloscopes, and radar displays. CRTs have been superseded by more modern display technologies such as LCD, plasma display, and OLED, which as of 2012 offer lower manufacturing and distribution costs. Television Computer Monitor

Robert Millikan Millikan’s Oil Drop Experiment
Starting in 1908, while a professor at the University of Chicago, Millikan worked on an oil-drop experiment in which he measured the charge on a single electron. J.J. Thomason had already discovered the charge-to-mass ratio of the electron. However, the actual charge and mass values were unknown. Therefore, if one of these two values were to be discovered, the other could easily be calculated. Millikan and his then graduate student Harvey Fletcher used the oil-drop experiment to measure the charge of the electron (as well as the electron mass, and Avogadro’s number, since their relation to the electron charge was known). Millikan showed that the results could be explained as integer multiples of a common value (1.592 × 10−19 coulomb), the charge on a single electron. That this is somewhat lower than the modern value of (14) x 10−19 coulomb is probably due to Millikan's use of an inaccurate value for the viscosity of air.

Charge and Mass of Electron- Robert Millikan
Charge of the electron is 1.592 × 10−19 coulomb The oil drop apparatus determined the actual e- charge: × 10−19 coulomb) (Somewhat lower than the modern of × 10 −19 coulomb) determined actual e- mass: 1/1837 the mass of H atom (9.11 x g) Robert Millikan's Oil drop experiment

Robert Millikan ( ) Doctorate in Physics from Columbia University 1923 Nobel Prize for Physics Determined charge on the electron as well as Planck’s constant Millikan set up his experiment by spraying oil drops, which fell between positively charged plate and negatively charged plate. He balanced the electric charge with electron's weight, as gravity pulled down on the oil drop, and the electric charges attracted and repelled it. Using this method, he discovered that the charge of an electron is negative. The weight of each drop was determined by observing it's rate of free fall through the air. Originally, he sprayed droplets of water through an electrical field that was intended to suspend them in midair, but the water droplets evaporated too fast, so Millikan switched to using oil.

Conclusions from Electron study:
Because atoms are electrically neutral, they must contain a positive charge to balance the negative e- Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass.

Discovery of Protons: Eugen Goldstein
Eugen Goldstein observed + particles (now called “proton”) particles with +charge & relative mass of 1 (or 1840x e-) called them “Canal rays” (anode rays)

Eugen Goldstein ( ) Credited by many for the discovery of the existence of protons discoverer of anode rays (canal rays) Worked in Berlin and Potsdam Observatories

Rutherford’s Gold Foil Experiment

Ernest Rutherford’s 1911- Gold Foil Experiment
Alpha particles fired at gold foil Particles hitting detecting screen were recorded

Discovery of Nucleus: Ernest Rutherford
most particles passed thru few deflected VERY FEW greatly deflected “It was as if you fired a 15-in [artillery] shell at a piece of tissue paper and it came back and hit you.” Conclusions: small, dense nucleus + charged nucleus

Rutherford’s Atomic Model
Experimental evidence: atom mostly empty space almost all mass in nucleus e- distributed around nucleus…occupy most of the volume called “nuclear model” Rutherford’s Atom 3:08

Discovery of Neutrons: James Chadwick
1932 – James Chadwick confirmed existence of the “neutron” particle with no charge n0 mass = p+ mass Nobel Prize for Physics Prepared the way for the fission of Uranium-235 which led to developing atomic bomb (Manhattan Project)

C. Atomic Scientists Democritus John Dalton J.J. Thomson
Robert A. Millikan With this information, scientists were able to determine the mass of an electron. named the atom proposed the Atomic Theory cathode-ray tube experiments measured the charge-to-mass ratio of an electron. (plum-pudding model) oil drop experiment measured the charge of an electron

C. Atomic Scientists Rutherford: gold foil experiment
James Chadwick: gold foil experiment Discovered small, dense nucleus Positive charged nucleus proved existence of neutrons.. no charge

D. Subatomic Particles 1- 9.11 x 10-28 1 1.67 x 10-24 1+ 1
Charge Mass Number Actual Mass (g) Location Electron (e-) Proton (p+) Neutron (no) Outside Nucleus (Electron cloud) 1- 9.11 x 10-28 Nucleus 1 1.67 x 10-24 1+ Nucleus 1 1.67 x 10-24

The Structure of the Atom
Chapter 3 Section 2 The Structure of the Atom The Structure of the Atom An is the smallest particle of an element that retains the chemical properties of that element. atom The is a very small region located at the center of an atom. nucleus The nucleus is made up of at least one positively charged particle called a and usually one or more neutral particles called . proton neutrons

Chapter 3 Section 2 The Structure of the Atom The Structure of the Atom Surrounding the nucleus is a region occupied by negatively charged particles called . Protons, neutrons, and electrons are often referred to as electrons subatomic particles

Chapter 3 Section 2 The Structure of the Atom The Structure of the Atom Atoms are electrically because they contain equal numbers of protons and electrons The nuclei of atoms of different elements differ in their number of and therefore in the amount of positive charge they possess. Thus, the number of determines that atom’s identity. neutral protons protons

Properties of Subatomic Particles
Chapter 3 Section 2 The Structure of the Atom Properties of Subatomic Particles

Forces in the Nucleus Chapter 3
Section 2 The Structure of the Atom Forces in the Nucleus When two protons are extremely close to each other, there is a strong attraction between them. A similar attraction exists when neutrons are very close to each other or when protons and neutrons are very close together. The short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together are referred to as nuclear forces.

Chapter 3 The Sizes of Atoms
Section 2 The Structure of the Atom The Sizes of Atoms The radius of an atom is the distance from the center of the nucleus to the outer portion of its electron cloud. Because atomic radii are so small, they are expressed using a unit that is more convenient for the sizes of atoms. This unit is the picometer, pm.

Vocabulary atomic number isotope mass number nuclide
unified atomic mass unit average atomic mass mole Avogadro’s number molar mass

Chapter 3 Objectives Explain what isotopes are.
Section 3 Counting Atoms Objectives Explain what isotopes are. Define atomic number and mass number, and describe how they apply to isotopes. Given the identity of a nuclide, determine its number of protons, neutrons, and electrons. Define mole, Avogadro’s number, and molar mass, and state how all three are related. Solve problems involving mass in grams, amount in moles, and number of atoms of an element.

Atomic Number Atoms of different elements have different numbers of
Atoms of the same element all have the same number of The (Z) of an element is the number of protons of each atom of that element. protons protons. atomic number

Atomic Number = number of protons

Atomic Number electrically neutral atoms: # of p+ = # of e-
Atomic number (Z) of element = # of protons in nucleus of each atom of that element. Element # of protons Atomic # (Z) Carbon 6 Phosphorus 15 Gold 79 # of electrons electrically neutral atoms: # of p+ = # of e-

Mass Number = # of protons + neutrons

X E. Nuclear Symbols Mass number Element Symbol Atomic number
Superscript → Element Symbol Atomic number Subscript →

Br Nuclear Symbols 80 35 Find each of these: number of protons
number of neutrons number of electrons Atomic number Mass Number 35 45 Br 80 35 35 35 80

Isotopes different mass #
Dalton wrong about elements of same type being identical… Element’s atoms can have different # neutrons. different mass # Isotopes (flavors)

Isotopes Frederick Soddy (1877-1956) proposed idea of isotopes in 1912
Isotopes - atoms of same element with different masses, because varying #s of neutrons 1921 Nobel Prize in Chemistry crater w/ his name on far side of Moon

Isotopes Elements occur in nature as mixtures of isotopes.
Although isotopes have different masses, they do not differ significantly in their chemical behavior

E. Hyphen notation We can also put mass number with a hyphen after name of element: carbon-12 carbon-14 uranium-235 uranium-238 Mass Number Mass Number

Hydrogen–1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2
Isotope Protons Electrons Neutrons Nucleus Hydrogen–1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2

Isotopes and Mass Number
Isotopes are atoms of the same element that have different The isotopes of a particular element all have the same number of protons and electrons but different numbers of Most of the elements consist of mixtures of isotopes. masses. neutrons. The number is the total number of and that make up the of an isotope. mass protons neutrons nucleus

Designating Isotopes uranium-235
Hyphen notation: The mass number is written with a hyphen after the name of the element. uranium-235 Nuclear symbol: The superscript indicates the mass number and the subscript indicates the atomic number.

uranium-235 Designating Isotopes
The number of neutrons is found by subtracting the atomic number from the mass number. uranium-235 mass # − atomic # = # of neutrons 235 (protons + neutrons) − 92 protons = 143 neutrons Nuclide is a general term for a specific isotope of an element.

Mass Number Mass # = p+ + n0 Isotope p+ n0 e- Mass # Oxygen - 10 - 33
Mass number is # of p+ and n0 in nucleus of an isotope: Mass # = p+ + n0 Isotope p+ n0 e- Mass # Oxygen - 10 - 33 42 - 31 15 18 8 8 18 Arsenic 75 33 75 Phosphorus 16 15 31

made up of 17 protons, 17 electrons.
Sample Problem A How many protons, electrons, and neutrons are there in an atom of chlorine-37? Atomic number = 17 made up of 17 protons, 17 electrons. Mass number = 37 mass number − atomic number = 37 (p+ plus n0) − 17 p+ = 20 n0 An atom of chlorine-37 is made up of 17 e-, 17 p+, and 20 n0.

Nuclear Symbols If an element has an atomic number of 34 and a mass number of 77, what is the: number of protons number of neutrons number of electrons hyphen notation nuclear symbol 34 43 34 selenium-77 Se 77 34

Nuclear Symbols If an element has 91 protons and 140 neutrons what is the Atomic number Mass number number of electrons hyphen notation nuclear symbol 91 231 91 protactinium-231 Pa 231 91

Nuclear Symbols If an element has 78 electrons and 117 neutrons what is the Atomic number Mass number number of protons hyphen symbol nuclear symbol 78 195 78 platinum-195 Pt 195 78

Chapter 3.3 Practice Problems
Set A p. 76 (1-3) 1. 2. 3. 35 protons 35 electrons 45 neutrons C 13 6 Phosophorus - 30

unified atomic mass unit:
Chapter 3 Section 3 Counting Atoms Relative Atomic Masses The standard used by scientists to compare units of atomic mass is the carbon-12 atom, which has been arbitrarily assigned a mass of exactly 12 atomic mass units, or 12 amu. unified atomic mass unit: One atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.

Measuring Atomic Mass Atomic Mass Unit (amu)
one-twelfth mass of C-12 atom Carbon-12 chosen b/c of its isotope purity Each isotope has own atomic mass we determine average from % abundance.

Average Atomic Mass How heavy is an oxygen atom?
Depends, b/c different kinds of oxygen atoms exist. We’re more concerned with average atomic mass. Based on abundance (%) of each variety of that element in nature. Don’t use grams - #’s tooooo small

Calculate the average:
Multiply atomic mass of each isotope by abundance (decimal), then add results. If not told otherwise, mass of isotope expressed in atomic mass units (amu)

Average Atomic Masses of Elements
Chapter 3 Section 3 Counting Atoms Average Atomic Masses of Elements Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. Calculating Average Atomic Mass The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes.

Chapter 3 Calculating Average Atomic Mass
Section 3 Counting Atoms Calculating Average Atomic Mass Copper consists of 69.15% copper-63, which has an atomic mass of amu, and 30.85% copper-65, which has an atomic mass of amu. The average atomic mass of copper can be calculated by multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results.

Chapter 3 Section 3 Counting Atoms ( × amu) + ( × amu) = amu The calculated average atomic mass of naturally occurring copper is amu.

Composition of the nucleus
Average Atomic Mass Atomic mass is average of all naturally occurring isotopes of that element. Atomic mass (amu) 12.00 13.00 14.00 Isotope Symbol Composition of the nucleus % in nature Carbon-12 C 6 protons 6 neutrons 98.89% Carbon-13 7 neutrons 1.11% Carbon-14 8 neutrons <0.01% 12 6 13 6 14 6 What is the average atomic mass of Carbon? 12.01

Relating Mass to Numbers of Atoms
Chapter 3 Section 3 Counting Atoms Relating Mass to Numbers of Atoms The Mole The mole is the SI unit for amount of substance. A mole (abbreviated mol) is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12. Avogadro’s Number Avogadro’s number— × 1023—is the number of particles in exactly one mole of a pure substance.

Chapter 3 Section 3 Counting Atoms Relating Mass to Numbers of Atoms Molar Mass The mass of one mole of a pure substance is called the of that substance. molar mass Molar mass is usually written in units of g/mol. The molar mass of an element is numerically equal to the atomic mass of the element in atomic mass units.

Chapter 3 Section 3 Counting Atoms Relating Mass to Numbers of Atoms Gram/Mole Conversions Chemists use molar mass as a conversion factor in chemical calculations. For example, the molar mass of helium is 4.00 g He/mol He. To find how many grams of helium there are in two moles of helium, multiply by the molar mass.

Chapter 3 Section 3 Counting Atoms Relating Mass to Numbers of Atoms Conversions with Avogadro’s Number Avogadro’s number can be used to find the number of atoms of an element from the amount in moles or to find the amount of an element in moles from the number of atoms. In these calculations, Avogadro’s number is expressed in units of atoms per mole.

Chapter 3 Section 3 Counting Atoms Solving Mole Problems

Chapter 3 Section 3 Counting Atoms

Chapter 3 Section 3 Counting Atoms Relating Mass to Numbers of Atoms Sample Problem B What is the mass in grams of 3.50 mol of the element copper, Cu? Given: 3.50 mol Cu Unknown: mass of Cu in grams Solution:

Chapter 3 Sample Problem B Solution
Section 3 Counting Atoms Sample Problem B Solution The molar mass of copper from the periodic table is rounded to g/mol.

Chapter 3 Sample Problem C
Section 3 Counting Atoms Sample Problem C A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced? Sample Problem C Solution Given: Unknown: Solution: 11.9 g Al amount of Al in moles

Chapter 3 Practice Problems
Set B p. 81 (1-4) * check answers in Appendix E Set C p. 81 (1-3) page R119

Chapter 3 Sample Problem D
Section 3 Counting Atoms Sample Problem D How many moles of silver, Ag, are in 3.01  1023 atoms of silver? Given: Unknown: Solution: 3.01 × 1023 atoms of Ag amount of Ag in moles

Chapter 3 Sample Problem E
Section 3 Counting Atoms Sample Problem E What is the mass in grams of 1.20  108 atoms of copper, Cu? Given: Unknown: Solution: 1.20 × 108 atoms of Cu mass of Cu in grams The molar mass of copper from the periodic table is rounded to g/mol.

Remember: Conversion Factors
Notecard Remember: Conversion Factors or or

Chapter 3 Practice Problems
Set D p. 82 (1-3) Set E p. 83 (1-3) * check answers in Appendix E page R119

End of Chapter 3 Show

Chapter 3 Atoms.

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