DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY Chemistry 140 Fall 2002 TENTH EDITION GENERAL CHEMISTRY Principles and Modern Applications PETRUCCI HERRING MADURA BISSONNETTE 2 Atoms and the Atomic Theory PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY

Atoms and the Atomic Theory Chemistry 140 Fall 2002 Contents 2-1 Early Chemical Discoveries and the Atomic Theory 2-2 Electrons and Other Discoveries in Atomic Physics 2-3 The Nuclear Atom 2-4 Chemical Elements 2-5 Atomic Mass 2-6 Introduction to the Periodic Table 2-7 The Concept of the Mole and the Avogadro Constant 2-8 Using the Mole Concept in Calculations Atoms and the Atomic Theory Image of silicon atoms that are only 78 pm apart; image produced by using a scanning transmission electron microscope (STEM). The hypothesis that all matter is made up of atoms has existed for more than 2000 years. It is only within the last few decades, however, that techniques have been developed that can render individual atoms visible.

2-1 Early Discoveries and the Atomic Theory Lavoisier 1774 Formulate the law of conservation of mass He heated a sealed glass vessel containing a sample of tin and some air The mass before heating (vessel + tin + air)=after heating (vessel + tin + air) The total mass of the substance present after a chemical reaction is the same as the total mass of substance before the reaction.

Figure show the reaction between silver nitrate and potassium chromate to give a red solid (silver chromate) (a) Before the reaction, the beaker with a silver nitrate solution and a graduated potassium chromate solution are placed on a single pan balance displace the combine mass = 104.5 g (b) After mixing, a chemical reaction occurs that forms silver chromate (red precipitate) in potassium nitrate solution. The total mass = 104.5 g, remains unchanged. FIGURE 2-2 Mass is conserved during a chemical reaction The low of conversation of mass says that matter is neither created nor destroyed in a chemical reaction

Sample A and Its Composition Proust 1799 Law of constant composition All sample of the compound have the same composition- the same proportions by mass of the constituent elements. Consider the compound water made up of two atoms of hydrogen (H) for every atoms of oxygen (O) Can be presented chemical formula H20 Two samples describes below have the same proportions of the two elements Exp: determine the percent by mass of hydrogen Simple divide the mass of hydrogen by the sample mass and multiply by 100%. For each sample, you will obtain the same results:11.19% H Sample A and Its Composition 10.000 g 27.000 g 1.119 g H % H=11.19 3.021 g H 8.881 g O % O= 88.81 23.979 g O

Dalton’s Atomic Theory: John Dalton Describes the basis of atomic theory with three assumptions Each element is composed of small particles called atoms. Atoms are neither created nor destroyed in chemical reactions. All atoms of a given element are identical but atoms of one element are different from those off all other elements Compounds are formed when atoms of more than one element combine in simple numerical ratios. exp: one atom of A to two B (AB2)

The second oxide is richer in O In forming carbon monoxide (CO), 1.0 g of carbon combines with 1.33 g of oxygen. In forming carbon dioxide (CO2), 1.0 g of carbon combines with 2.66 g of oxygen. Figure 2-3 Molecules of CO and CO2 The second oxide is richer in O It contains twice as much O as the first

2-2 Electrons and Other Discoveries in Atomic Physics Electricity and magnetism were used in the experiment so that led to the current theory of atomic structure Certain objects displays a properties called electric charge, which can be either positive (+) or negative (-) An object having equal number of (+) or (-) charged particles carries no net charge and is electrically neutral If the number of (+) charge exceed the number of (-) charge , the object has a net positive charge If the number of (-) charge exceed the number of (+) charge , the object has a net negative charge

Forces between electrically charged objects (+) and (-) charges attract each other , while two (+) and two (-) charges repel each other FIGURE 2-4 Forces between electrically charged objects (a) Electrostatically charged comb. If you comb your hair the static charge develop on the comb and causes bits of paper to be attracted to the comb (b) Both objects on the left carry negative charge repel each other The objects in the center lack any electrical charge and exert no force on each other The object on the right carry opposite charges and attract each other

The Discovery of Electrons Faraday discovered cathode rays, a type of radiation emitted by a (-) terminal, or cathode (it is iron, platinum so on) The radiation crossed the evacuated tube to the (+) terminal, or anode FIGURE 2-6 Cathode ray tube The high voltage source of electricity creates a (-) charge on the electrode at the left (cathode) and a (+) charge on the electrode at the right (anode) Cathode rays pass from the cathode (C) to the anode (A) which is perforated to allow the passage of a narrow beam of the cathode rays They are visible only through the green florescence that they produce on the zinc sulfide-coated screen at the end of the tube. They are in the other part s of the tube

Cathode rays and their properties C rays are deflected by electric and magnetic fields in the manner expected for negatively charged particles (Figure 2-7 (a) and (b)) Figure 2-7 (c) determine the mass to charge ratio m/e for the C rays Cathode rays subsequently become known as electrons Electron m/e = -5.6857 × 10-9 g coulomb-1 FIGURE 2-7 Cathode rays and their properties

Millikan’s oil-drop experiment Robert Millikan determined the electronic charge , e through a serious oil drop experiment He showed ionized oil drops can be balanced against the pull of gravity by an electric field e is -1.6022x10-19 coulomb mass of electron= (-1.6022x10-19 coulomb) x (-5.6857 × 10-9 g coulomb-1 ) = 9.1094 x 10-28 g Figure 2-8 Millikan’s oil-drop experiment

Plum-pudding Model Proposed by Thomson Explains how the electron particles were incorporated into atoms. He thought that the positive charge necessary to counter balance the negative charges of electrons in a neutral atom was the form of a nebulous cloud. He suggest, electrons floated in a diffuse cloud of positive charge A helium atom would have a +2 cloud of (+) charge and two electrons (-2) If helium atom loses one electron, it becomes charged and is called an ion referred to He+ has a net charge of 1+ If the helium atom loses both electron the He2+ ion forms

X-Rays and Radioactivity X-ray is form of high energy electromagnetic radiation Radioactivity is the spontaneous emission of radiation from a substance Two types of radiation form from radioactive material were identified by Ernest Rutherford Alpha (a): a-particles carry two fundamental units of positive charge and the same mass as helium atoms. This particle are identical to He2+ions Beta (b): b-particles are negatively charged and have the same properties as electrons Gamma (g) rays: is not effected by electric or magnetic field. It is not made of particles. It is electromagnetic radiation of extremely high penetrating power.

2-3 The Nuclear Atom The scattering of a- particles by metal foil Rutherford used the a-particles to study inner structure of the atoms The telescope travels in a circular track around at evacuated chamber containing the metal foil. Most a-particles pass thought the metal foil undeflected , but some are deflected through large angles Figure 2-11 The scattering of a- particles by metal foil The nuclear atom have these features below Most of mass and all of positive charge of an atom are centered in a very small region called nucleus. The remainder of the atom is mostly empty space The magnitude of the positive charge is different for the different atoms and is approximately one-half the atomic weight of the element There are as many electrons outside the nucleus as there are unit of positive charge on the nucleus. The atom as a whole is electrically neutral.

The nuclear atom – illustrated by the helium atom Properties of Protons, neutrons and Electrons Protons: positively charged fundamental particles of the matter in the nuclei of atoms Neutrons: penetrating radiation consisted of beam of neutral particles The number of protons in a given atom is called the atomic number, or the proton number, Z The number of electrons in the atom is equal to Z because the atom is electrically neutral The total number of proton and neutrons in an atom is called the mass number, A The number of neutron is A-Z and electrically neutral. Figure 2-13 The nuclear atom – illustrated by the helium atom

To represent a particular atom we use symbolism Chemistry 140 Fall 2002 2-4 Chemical Elements Each element has a name and distinctive symbol Exp: carbon:C, oxygen:O, neon:Ne, iron:Fe To represent a particular atom we use symbolism Symbol of element A= mass number Z = atomic number Often do not specify Z when writing. For example 14C, C specifies Z = 12. Special names for some isotopes. For example hydrogen, deuterium, tritium. Has 13 protons and 14 neutrons in its nucleus and 13 electron outside the nucleus (recall that an atom has the same number of electrons as protons) Al 13 27

Chemistry 140 Fall 2002 Isotopes atoms that have the same atomic number (Z) but different masss number (A) are called isotopes. Exp: all neon atoms have 10 protons in their nuclei, and most have 10 neutron as well. A very few neon atoms have 11 neutrons and some have 12 Ne 10 20 Ne 10 21 Ne 10 22 Ions When atoms lose or gain electrons the species formed are called ions and carry net charges. Removing electrons result in positively charged ion The number of proton does not change when an atom becomes an ion. Exp: Ne 10 20 + 10 protons 10 neutrons and 9 electrons Often do not specify Z when writing. For example 14C, C specifies Z = 12. Special names for some isotopes. For example hydrogen, deuterium, tritium. Ne 10 22 2+ 10 protons 12 neutrons and 8 electrons O 8 16 2- 8 protons 8 neutrons and 10 electrons

A mass spectrometer and a mass spectrum Isotopic masses Used when original mass of an atoms can not be determined That must be done by experiment One type of atom has been chosen and assigned a specific mass. This standard is an atom of the isotope carbon-12 Next the masses of the other atoms relative to carbon -12 are determined with a mass spectrometer Figure 2-14 A mass spectrometer and a mass spectrum

2-5 Atomic Mass The average of the isotopic masses, weighted according to the naturally occurring abundance of the isotopes of the elements Weighted Average Atomic Mass of an Element Equation (2.3) fractional abundance of isotope 1 atomic mass of isotope 1 fractional abundance of isotope 2 atomic mass of isotope 2 x + = x + …… Aave = x1 x A1 + x2 x A2 + …… xn x An where x1 + x2+ …..+ xn = 1.0

2-6 Introduction to The Periodic Table The classification system we need known as the periodic table of the elements Read atomic masses Read the ions formed by main group elements Read the electron configuration Learn trends in physical and chemical properties We will discuss these in detail in Chapter 9.

The Periodic table Noble Gases Alkali Metals Main Group Alkaline Earths Halogens Transition Metals Lanthanides and Actinides

2-7 The Concept of the Mole and the Avogadro Constant A mole: is the amount of the substance that contains the same number of elementary entities (atoms, molecules and so on) Avogadro constant or Avogadro number, NA: The amount of elementary entities in a mole NA = 6.02214179 x 1023 mol-1 Exp: 1 mol 12C = 6.02214179 x 1023 12C atoms = 12 g 1 mol 16O = 6.02214179 x 1023 16O atoms = 15.9949 g (and so on) Molar mass, M: the mass of one mole of substance , from a table of atomic masses Exp: the molar mass of lithium is 6.941 g/mol Li