The Structure of an Atom Chapter 4
4.1: Defining the Atom Essential Question: How did the concept of the atom change from the time of Democritus to the time of John Dalton?
4.1: Defining the Atom All matter is composed of particles, which are called atoms. Atoms are the smallest particle of an element that retains its identity in a chemical reaction.
4.1: Defining the Atom The Greek philosopher Democritus (460 B.C.-370 B.C) was among the first to suggest the existence of atoms. Democritus reasoned that atoms were indivisible and indestructible.
4.1: Defining the Atom Democritus’s ideas agreed with later scientific theory. His ideas did NOT explain chemical behavior. His ideas lacked experimental support because his approach was not based on the scientific method. (It was based in philosophy and logic)
4.1: Defining the Atom The modern process of discovery regarding atoms began with John Dalton (1766-1844), an English chemist. By using experimental methods, Dalton transferred Democritus’s ideas on atoms into scientific theory.
4.1: Defining the Atom John Dalton’s theory stated that: 1. All elements are made up of tiny indivisible particles called atoms. 2. Atoms of the same element are identical. 3. Atoms of different elements can physically mix together OR can chemically combine to form compounds in simple whole # ratios. 4. Chemical reactions occur when atoms are separated and rearrange.
How small is an atom? Atoms are very small. The radii of most atoms fall within a range of 5 x 10-11m to 2 x 10-10m Observable with scanning electron microscopes. Electron microscopes are capable of higher magnification than light microscopes.
Scanning Electron Microscope
4.2: Structure of the Nuclear Atom Subatomic Particles: Most of Dalton’s theory is acceptable today. However, one important change is that atoms are now known to be divisible. They can be broken down into even smaller, most fundamental particles, called subatomic particles. There are 3 types: 1. Electrons 2. Protons 3. Neutrons
4.2: Structure of the Nuclear Atom J. J. Thompson (1858-1940) 1897: Discovered the electron. Electrons are negatively charged subatomic particles.
4.2: Structure of the Nuclear Atom Thomson performed experiments that involved passing electric current through gases with low pressure. He sealed the gases in the glass tubes fitted at both ends with electrodes. One electrode, the anode, became positively charged. The other electrode, the cathode became negatively charged. The result was a glowing beam, or cathode ray that traveled from the cathode to the anode. Show video clip here
4.2: Structure of the Nuclear Atom Protons and Neutrons: In 1886, Eugene Goldstein observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. He called those rays, canal rays and concluded that they were composed on positive particles called protons. Protons are positively charged subatomic particles.
4.2: Structure of the Nuclear Atom Protons and Neutrons: In 1932, James Chadwick confirmed the existence of the neutron. Neutrons are subatomic particles with no charge.
4.2: Structure of the Nuclear Atom The Atomic Nucleus: Rutherford’s Gold Foil Experiment: Physicist Ernest Rutherford established the nuclear theory of the atom with his gold-foil experiment. When he shot a beam of alpha particles as a sheet of gold foil, a few of the particles were deflected. He concluded that a tiny, dense nucleus was causing the deflections.
Ernest Rutherford Discovered the nucleus of an atom
Matter is mostly empty space
4.2: Structure of the Nuclear Atom The Atomic Nucleus: The Rutherford Atomic Model: Rutherford suggested a new theory of the atom. He proposed that the atom is mostly empty space, thus explaining the lack of deflection of most of the alpha particles. He concluded that all of the positive charge and almost all of the mas are concentrated in a small region that has enough positive charge to account for the great deflection of some of the alpha particles.
4.2: Structure of the Nuclear Atom The Atomic Nucleus: The Rutherford Atomic Model: The nucleus is the tiny central core of an atom and is composed of protons and neutrons.
4.2: Structure of the Nuclear Atom The Atomic Nucleus: The Rutherford Atomic Model: The Rutherford atomic model is known as the nuclear atom. In the nuclear atom, the protons and neutrons are located in the positively charged nucleus. The electrons are distributed around the nucleus and occupy almost all of the volume of the atom. Do Rutherford Scattering Activity Here
End Lesson #1
4.3: Distinguishing Among Atoms Atomic Number and Mass Number: Atoms are composed of protons, neutrons and electrons. Protons and neutrons make up the nucleus and electrons surround the nucleus. Elements are different because they contain different numbers of protons. An element’s atomic number is the number of protons in the nucleus.
} A. Basic Structure + + proton (+) neutron (Ø) Nucleus electron (-) electron cloud neutron (Ø) } + + Nucleus Nucleus: smallest yet heaviest part of the atom
4.3: Distinguishing Among Atoms Mass Number: The total number of protons and neutrons in an atom is called the mass number. Number of neutrons = Mass Number – Atomic number
4.3: Distinguishing Among Atoms Key Point: Elements are different because they contain different numbers of protons. An element’s atomic number is the number of protons in the nucleus of an atom of that element. Mass number or atomic mass: The total number of protons + neutrons in an atom.
Hydrogen 1 H 1.008 Element Name Atomic number Element Symbol Avg. Atomic Mass
End Lesson #2
4.3: Distinguishing Among Atoms EQ: How do isotopes of an element differ? 1. Isotopes: Isotopes are: atoms of the same element that have different numbers of neutrons. (which means they also have different mass numbers) Isotopes are chemically alike because they have the same number of protons.
4.3: Distinguishing Among Atoms 2. Analogy: Isotopes are like Skittles – Skittles have different “flavors” just like atoms of the same element have different “flavors” or isotopes.
4.3: Distinguishing Among Atoms
4.3: Distinguishing Among Atoms 3. Isotopic Notation Mass Notation Charge Lithium - 7 Mass # → 7 + Li mass # element name 3 Atomic # →
4.3: Distinguishing Among Atoms 4. Average Atomic Mass is: Weighted average of all isotopes The isotope that is the most abundant has the greatest effect on the avg. atomic mass found on the periodic table. Example: The average atomic mass of copper is 63.546 amu. Which of copper’s two isotopes is more abundant: Copper-63 or Copper-65? Why?
4.3: Distinguishing Among Atoms 3. Avg. Atomic mass = ∑ (mass x abundance)
4.3: Distinguishing Among Atoms Avg. Atomic mass = ∑ (mass x abundance) Example: Boron has 2 naturally occurring isotopes. Boron-10 has a mass of 10.012amu and a relative abundance of 19.91%. Boron-11 has a mass of 11.009amu and a relative abundance of 80.09%. What is the average atomic mass of Boron?
4.3: Distinguishing Among Atoms Example: Boron has 2 naturally occurring isotopes. Boron-10 has a mass of 10.012amu and a relative abundance of 19.91%. Boron-11 has a mass of 11.009amu and a relative abundance of 80.09%. What is the average atomic mass of Boron? Mass x Abundance = Atomic Mass B-10 = 10.012amu x 0.1991 = B-11 = 11.009amu x 0.8009 = + Avg. Atomic mass
End Lesson #3
Robert Millikan (1916) Discovered the quantity of charge of an electron (1.60 x 10-19 coulomb) An electron has one unit of negative charge. An electron’s mass is 1/1840th mass of one hydrogen atom (actual mass = 9.11 x 10-28)
Hydrogen 1 H 1.008 Element Name Element Symbol Avg. Atomic Mass Atomic number # of protons Element Symbol Avg. Atomic Mass
EQ: How are ions and atoms different? Ion Notes EQ: How are ions and atoms different?
A. Definitions Ion: When an atom gains or loses an electron, it has an imbalance of charge Cation (+) = loses e- to become positive Anion (-) = gains e- to become negative In an ion, the number of protons do not equal the number of electrons
B. Examples Element Atomic # Mass # Protons Neutrons Electrons O2- F- Ca2+ 16 8 8 16 8 8 10 8 19 9 19 9 10 10 9 41 20 41 20 21 18 20
Unstable Nuclei
A. Normal Reactions atoms rearrange, the elements do not change
B. Nuclear Reactions 14 14 C N + 6 7 -1 Radioactive decay- atom breaks apart spontaneously 14 14 C N + 6 7 -1 *Note: Please fit all reactions in one line
B. Nuclear Reactions 9 4 12 1 Be n He C + + 4 6 2 Radioactive bombardment: Particle hits atom & it splits 9 4 12 1 Be n He C + + 4 6 2
C. Types of Radioactive Particles Symbol Composition Penetrating Power Alpha, He 2 P & 2 N Low Beta, electron 100 x alpha gamma, 0 EM waves Very great 4 2 -1
Penetrating Power Alpha Beta Gamma http://www.fusrapmaywood.com/factsheet/..%5Cimages%5Cfoil2.gif
Penetrating Power http://www.bcm.edu/bodycomplab/Images/pntrtn.gif
D. Miscellaneous Notation 1. Positron 2. Neutron e +1 1 n
E. Transmutation 1. Fission : a very heavy-mass nucleus splits to form two medium-mass (size) nuclei.
E. Transmutation 2. Fusion : two very light-mass nuclei combine to form heavier, more stable nuclei
F. Balancing Nuclear Equations 1. mass # & atomic #’s must add up the same on both sides of the equation 31 4 27 1 30 Al + He + H Si ____ 2 13 1 14 15
Example #1 14 14 C N ? + 6 7 14 + ? = 14 7 + ? = 6 -1 e -1
Example #2 230 4 Th ? He + + 2 90 ? + 4 + 0 = 230 226 88 ? + 2 + 0 = 90 226 Ra 88
Example #3 4 27 30 Al + He Si + ? 2 13 14 27 + 4 = 30 + ? 1 1 13 + 2 = 14 + ? 1 1 p or H 1 1