Chapter 4: Atomic Structure By Kendon Smith – Columbia Central HS – Brooklyn, MI
I. Early Models of the Atom - All matter is composed of tiny particles called atoms, which are the smallest particles of an element that retain its properties and identity during a chemical reaction.A. Democritus’s Atomic Philosophy- Democritus was a Greek philosopher from the 4th century B.C. who first suggested the existence of tiny particles called “atomos”.- Democritus believed atoms were indivisible and indestructible.- Democritus lacked experimental support for his ideas.
B. Dalton’s Atomic Theory- John Dalton (1766 – 1844) was an English chemist and schoolteacher.- Dalton used experiments to transform Democritus’ ideas into scientific theory.
Dalton’s Atomic Theory:1. All elements are composed of tiny indivisible particles called atoms.2. Atoms of the same element are identical. Atoms of one element are different than atoms of another element.3. Atoms of different elements can physically mix together, or they can chemically combine in simple whole-number ratios to form compounds.4. Chemical reactions occur when atoms are separated, joined, or rearranged, however atoms of one element are never changed into atoms of a different element.
C. Sizing up the Atom- Atoms are so tiny that a single copper penny contains 2.4 x 1022 atoms. - A line of 100,000,000 copper atoms would be 1 centimeter long. - The radii of most atoms are between 5 x 10-11 m and 2 x 10-10 m. Calculate the diameter range of most atoms in pm:Radius = 5 x 10-11 m Radius = 2 x 10-10 mDiameter = 1 x 10-10 m Diameter = 4 x 10-10 m ÷ 10-12 = 100 pm ÷ 10-12 = 400 pm
II. Structure of the AtomA. Subatomic Particles- Atoms are now known to be divisible. They can be broken down into even smaller particles, called subatomic particles.- The three subatomic particles are electrons, protons, and neutrons.
1. Electrons a. J. J. Thomson, an English physicist, discovered the electron in 1897. b. Electrons are negatively charged subatomic particles. c. Thomson performed the cathode ray tube experiment, in which a beam of negatively charged particles traveled from the negative electrode, called the cathode, to the positive electrode, called the anode.
d. The ray was deflected by magnets and charged metal plates. It was repelled by a negative plate and attracted by a positive plate. Because Thomson knew that opposites attract, electrons must be negative.e. U.S. physicist Robert Millikan carried out experiments to measure the mass and charge of the electron.- An electron carries exactly one unit of negative charge = -1.- An electron’s mass is 1/1840 the mass of a proton ≈ basically ZERO mass.
2. Protons and Neutronsa. The cathode ray tube experiment taught us some simple concepts about atoms:1. Atoms have no net charge; they are electrically neutral.2. Electric charges are carried by particles of matter.3. Electric charges always exist in whole numbers – no fractions of charge.4. When equal numbers of negatives and positives join, particles are neutral.
b. This meant there must be a positive particle left behind when atoms lose their negative charged electrons!c. In 1886, Eugen Goldstein discovered positive particles called protons.d. In 1932, English physicist James Chadwick discovered neutrons.- Neutrons carry no charge and have a mass nearly equal to a proton.- Neutrons only contribute mass to an atom, making some atoms heavier.
RelativeParticle Symbol Charge Mass 1/1840 =Electron e - -1 zero!Proton p+ +1 1Neutron n0 0 1
B. The Atomic Nucleus 1. Rutherford’s Gold Foil Experimenta. In 1911, Ernest Rutherford tested the current atomic theory by shooting alpha particles at a very thin sheet of gold foil. - Alpha particles are helium atoms that have lost their electrons. They are made of two protons and two neutrons, so they have a double positive charge = +2.b. It was expected that the alpha particles would pass through the gold foil but experience some
c. Surprisingly, a majority of the alpha particles passed through the gold foil as if there was nothing there, with a few even bouncing back!d. This led Rutherford to two important conclusions about atoms: 1. Atoms are mostly empty space! (Explains lack of deflections.) 2. All the positive charge and mass of the atom must be located in a tiny, dense region in the central core of the atom, called the nucleus. (Explains occasional bounce backs.)e. In the nuclear atom, protons and neutrons are located in the nucleus. The electrons are distributed in the space around the nucleus.
C. The Bohr Model (from Chapter 5)1. Niels Bohr improved on Rutherford’s model of the atom and proposed that the electrons travel around the nucleus in specific circular paths, or orbits.- Orbitals, or energy levels, are larger as you move further away from the nucleus and can hold more electrons. Level 1 = 2 e- Level 2 = 8 e- Level 3 = 18 e- Level 4 = 32 e-
Models of the Atom: Thomson’s “Plum Pudding” ModelElectron (–) Positive Matrix (+)
Models of the Atom: Electron (–) Cloud Nucleus (+)(Protons & Neutrons)
Models of the Atom: The Bohr Model of Electron OrbitalsElectron (–)Energy Levelsor OrbitalsNucleusProtons (+) ++ + +Neutrons (0)
III. Distinguishing Among AtomsA. Atomic Number- Elements are different because they contain different numbers of protons.a. Atomic Number = the number of protons in the nucleus of an atom
B. Mass Numbera. Mass Number = the total number of protons and neutrons in the nucleusb. Only protons and neutrons add mass to an atom – Electrons are negligible! * Mass number is NOT THE SAME as Atomic Mass!c. Mass numbers are always whole numbers!d. The number of neutrons is the difference between mass number and atomic number. # of neutrons = mass number – atomic number
Atomic Number (protons) Atomic Mass (not mass #)
C. Isotopes- Atoms of the same element can have different numbers of neutrons, which gives them different mass numbers! a. Isotopes are atoms with the same number of protons, but different number of neutrons. b. Isotopes are atoms with the same atomic number, but different mass numbers.
IV. Atomic MassA. Atomic Mass Units (amu)- Even the largest atom is incredibly small!- A proton has an actual mass of 1.67 x 10-24 grams, so it is difficult to work with numbers these small.- Atomic mass units are units of relative mass that were invented to make the numbers easier to work with and understand.- Atomic mass units are based on the mass of an atom of the isotope Carbon-12, which has a mass of exactly 12 amu’s.
IV. Atomic MassA. Atomic Mass Units (amu)- 1 atom Carbon-12 = 12 amu’s- Therefore, 1 amu = 1/12th the mass of Carbon-12- Carbon-12 has 6 protons and 6 neutrons = 12 total particles in the nucleus- Therefore the mass of 1 p+ or n0 = 1 amu- What does relative mass tell us? It does not tell us the actual mass of an atom, but instead it tells us how it’s mass compares to the standard, which is Carbon-12.
Element Relative Mass MeaningMagnesium 24 amu 1 atom of Mg is 2x heavier than an atom of C-12 Helium 4 amu He is 3 times lighter than CTitanium 48 amu Ti is 4 times heavier than C
B. Calculating Atomic Mass Values- In nature, most elements occur as a mixture of isotopes- Each isotope has a different mass number, so the value used to describe the mass of these mixed samples is a type of average.- Average masses are weighted according to percent abundance, which means that those isotopes that are more abundant have a greater influence on the average mass.
NEED TO KNOW:a. How many isotopes exist for an element?b. What are the mass numbers of each isotope?c. What is the percent abundance for each isotope?Calculation Steps:1. Multiply each mass number by its % abundance. (% must be re-written as a decimal!)2. Add up all the results for the total weighted average atomic mass.
Sample Problems:14. Boron has two isotopes: boron-10 and boron-11. Which is more abundant, given that the atomic mass of boron is 10.81 amu? 10.81 BORON-10 BORON-11 10.5 Straight average? = ________ 10.81 Weighted average? = ________ (closer to 11!)
Sample Problems:15. There are three isotopes of silicon; they have mass numbers of 28, 29, and 30. The atomic mass of silicon is 28.086 amu. 28 29 30 Weighted Average = 28.086 amu
Sample Problems:16. The element copper has naturally occuring isotopes with mass numbers of 63 and 65. The relative abundance values are 69.2% for 63 amu, and 30.8% for 65 amu. Calculate the average atomic mass of copper. 63 amu x 0.692 = 43.596 ADD THEM 65 amu x 0.308 = 20.02 UP! 63.616 amu