Atomic Structure

Dalton’s Atomic Theory John Dalton (1766-1844) had four theories All elements are composed of submicroscopic indivisible particles called atoms Atoms of the same element are identical. The atoms of anyone element are different from those of any other element Atoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compounds Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction

Atom Atom- smallest particle of an element that retains the properties of that element Dalton’s Model of the Atom: Billiard Ball Model

Basic Laws Conservation of Mass Law of Definite Proportion- compounds have a constant composition by mass. They react in specific ratios by mass. Multiple Proportions- When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.

What?! Water has 8 g of oxygen per 1 g of hydrogen. Hydrogen peroxide has 16 g of oxygen per 1 g of hydrogen. 16/8 = 2/1 Small whole number ratios

Example of Law Of Multiple Proportions Mercury has two oxides. One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass. Show that these compounds follow the law of multiple proportion. Speculate on the formula of the two oxides.

Assume 100 g sample 96.2/3.8 = 25.3 g Hg/1 g O 25.3/12.5= 2/1of Hg Hg2O and HgO Amount Hg Amount O Compound A 96.2 g 3.8 g Compound B 92.6g 7.4 g

A Helpful Observation Gay-Lussac- under the same conditions of temperature and pressure, determined that gas compounds always react in whole number ratios by volume. Avogadro- interpreted that to mean at the same temperature and pressure, equal volumes of gas contain the same number of particles (called Avogadro’s hypothesis)

Dalton’s Atomic Theory Revised after New Evidence Part 2: Discovery of Isotopes: Atoms of the same element can have different masses Part 3: Atoms are Divisible made up of subatomic particles Part 3: Atoms cannot be created nor destroyed in ordinary chemical reactions They can in NUCLEAR REACTIONS

Electron J.J Thomson (1856-1940) – discovered the electron in 1897 Electron is the negative charged subatomic particle An electron carries exactly one unit of negative charge & its mass is 1/1840 the mass of a hydrogen atom

Cathode Ray The Cathode Ray tubes pass electricity through a gas that is contained at a very low pressure

The Electron Figure 2.4 Streams of negatively charged particles were found to emanate from cathode tubes. © 2009, Prentice-Hall, Inc.

The Electron Figure 2.4 Thompson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/g. © 2009, Prentice-Hall, Inc.

Thomson’s Atomic Model Plum Pudding Model Thomson thought electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model

Mass of Electron In 1909 Robert Millikan determined the mass of an electron with his Oil Drop Experiment He determined the mass to be 9.109 x 10-31 kg The oil drop apparatus

Millikan’s Experiment Atomizer - + Oil Microscope

Millikan’s Experiment Atomizer Oil droplets - + Oil Microscope

Millikan’s Experiment X-rays X-rays give some drops a charge by knocking off electrons

Millikan’s Experiment +

Millikan’s Experiment - - + + They put an electric charge on the plates

Millikan’s Experiment - - + + Some drops would hover

Millikan’s Experiment - - - - - - - + + + + + + + +

Millikan’s Experiment - - + + Measure the drop and find volume from 4/3πr3 Find mass from M = D x V

Millikan’s Experiment - - + + From the mass of the drop and the charge on the plates, he calculated the charge on an electron

Radioactivity Discovered by accident Bequerel Three types alpha- helium nucleus (+2 charge, large mass) beta- high speed electron gamma- high energy light

Ernest Rutherford Rutherford (1871-1937) proposed that all mass and all positive charges are in a small concentrated region at the center of the atom He used the Gold-Foil Experiment to prove his theory & in 1911 he discovered the nucleus Nucleus- central core of an atom, composed of protons and neutrons The nucleus is a positively charged region and it is surrounded by electrons which occupy most of the volume of the atom

Rutherford’s Experiment Used uranium to produce alpha particles Aimed alpha particles at gold foil by drilling hole in lead block Since the mass is evenly distributed in gold atoms alpha particles should go straight through. Used gold foil because it could be made atoms thin

Florescent Screen Lead block Uranium Gold Foil

What he expected

Because

Because, he thought the mass was evenly distributed in the atom

What he got

How he explained it Atom is mostly empty space Small dense, positive piece at center Alpha particles are deflected by it if they get close enough +

+

Modern View The atom is mostly empty space Two regions Nucleus- protons and neutrons Electron cloud- region where you have a chance of finding an electron

Neutron James Chadwick (1891-1974) – discovered the neutron in 1932 Neutron is a subatomic particle with no charge but their mass is nearly equal to that of a proton

Quark Protons & Neutrons can still be broken down into a smaller particle called the Quark The Quark is held together by Gluons

Subatomic particles Relative mass Actual mass (g) Name Symbol Charge Electron e- -1 1/1840 9.11 x 10-28 Proton p+ +1 1 1.67 x 10-24 Neutron n0 1 1.67 x 10-24

Sub-atomic Particles Z - atomic number = number of protons determines type of atom A - mass number = number of protons + neutrons in the nucleus Number of protons = number of electrons if neutral

Rules for Atomic Structure Atomic # = # of Protons # of Protons = # of Electrons if neutral Protons – Electrons = charge if not neutral Mass # = # of Protons + # of Neutrons M.A.N : mass number – atomic number = # neutrons

Isotopes Isotope- atoms that have the same number of protons but different number of neutrons Since isotopes have a different number of neutrons the isotope has a different mass number. Isotopes are still chemically alike because they have the same number of protons and electrons

Examples of Isotopes

Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C © 2009, Prentice-Hall, Inc.

X X Na Nuclear Symbols Mass number Atomic number Contain the symbol of the element, the mass number and the atomic number X Mass number X A Z Atomic number 23 Na 11

Nuclear Symbols Proton = p+ Electron = e- Neutron = n0 Charge Mass # - Superscript Atomic # - Subscript

Naming Isotopes Put the mass number after the name of the element carbon- 12 carbon -14 uranium-235

Na Electrical Charges Neutral atoms have no net electrical charges Given the number of negative charges combines with the number of positive charges = Electrically Neutral 11 protons 10 electrons 12 neutrons 23 1+ Na 11

Example 1 Write the isotopic symbols for three isotopes of Silicon in which there are 14, 15, and 16 neutrons 28 Si 14 29 Si 14 30 Si 14

Au Example 2 197 79 Find the number of protons number of neutrons number of electrons Atomic number Mass Number 79 118 79 79 197

Example 3 How many neutrons are in the 3 isotopes of magnesium-24, magnesium-25, and magnesium-26? 24 Mg 12 12 n 25 Mg 13 n 12 26 Mg 14 n 12

Example 4 Write the nuclear symbol for the atom that has an atomic number of 9 and a mass number of 19. How many electrons and neutrons does this atom have? 19 F 9 9 electrons 10 neutrons

Example 5 Write the formula for the fluoride ion. Its mass number is still 19. How many protons, neutrons and electrons are there? 1- 19 F 9 9 protons 10 electrons 10 neutrons

Example 6 How many protons, neutrons & electrons are found in the calcium-40 ion? 2+ 40 Ca 20 20 protons 18 electrons 20 neutrons

Measuring Atomic Mass Unit is the Atomic Mass Unit (amu) One twelfth the mass of a carbon-12 atom Each isotope has its own atomic mass. We need the average from the percent abundance Each isotope of an element has fixed mass and a natural % abundance You need both of these values to find the Atomic Weight

Calculating Atomic Weight Cl-35 34.969amu and 75.77% abundance Cl-37 36.966amu and 24.23% abundance To solve for Cl-35 AMU x Abundance 34.969 x .7577 = 26.496 You solve for Cl-37

Atomic Weight Cont. Cl-37 AMU x Abundance 36.966 x .2423 = 8.957 Now you combine your two answers 26.496 + 8.957= 35.453 Look at Cl on the table. What is the Atomic Weight?

Example 7 There are 2 isotopes of lithium, one with a mass of 6.015 amu and an abundance of 7.42%. The other isotope has a mass of 7.016 amu and an abundance of 92.58%. Calculate the atomic weight (average atomic mass) of lithium.

Example 8 Calculate the atomic weight of copper. Copper has two isotopes. One has 69.1% and has a mass of 62.93 amu. The other has a mass of 64.93 amu. What is the atomic weight???

Example 9 There are two isotopes of nitrogen, one with an atomic mass of 14.0031amu and one with a mass of 15.0001 amu. What is the percent abundance of each? Nitrogen’s average atomic mass is 14.0067 amu. X+ Y = 1 so 1st isotope = X 2nd isotope (Y) = 1-X

Atomic Weight & Decimals Atomic Weight- of an element is a weighted average mass of the atoms in a naturally occurring sample of an element Atomic Weights use decimal points because it is an average of an element

Experimental Determination of Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer. Figure 2.13 © 2009, Prentice-Hall, Inc.