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GENERAL CHEMISTRY. Principles and Modern Applications. TENTH EDITION. 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.

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Tenth edition

GENERAL CHEMISTRY

Principles and Modern Applications

TENTH EDITION

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

Atoms and the Atomic Theory


2 1 early discoveries and the atomic theory

2-1 Early Discoveries and the Atomic Theory

  • Lavoisier 1774Formulate 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 substancepresent after a chemical reaction is the same as the total mass of substance before the reaction.


Tenth edition

  • 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


Tenth edition

  • Proust 1799Law 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


Dalton s atomic theory john dalton

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)


Molecules of co and co 2

  • 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.

  • The second oxide is richer in O

  • It contains twice as much O as the first

Figure 2-3

Molecules of CO and CO2


2 2 electrons and other discoveries in atomic physics

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

Forces between electrically charged objects

  • (+) and (-) charges attract each other , while two (+) and two (-) charges repel each other

FIGURE 2-4

  • (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


Cathode ray tube

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

  • 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 thecathode rays

  • They are visible only through the green florescence that they produce on the zinc sulfide-coated screen at the endof the tube. They are in the other part s of the tube

FIGURE 2-6

Cathode ray 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

Plum-pudding Model Proposed by Thomson

  • Explains how the electron particles were incorporated into atoms.

  • He thought that thepositive charge necessary to counter balance the negative charges of electrons in a neutral atom wasthe 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-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 unitsof positive charge and the same mass as helium atoms. This particle areidentical 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 radiationof extremely high penetrating power.


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

  • 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. Theremainder 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.

Figure 2-11

The scattering of a-particles by metal foil

2-3 The Nuclear Atom


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


2 4 chemical elements

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 numberZ = atomic number

27

Al

13

  • 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)


Isotopes

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

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:

16

20

20

21

22

22

Ne

Ne

Ne

Ne

O

Ne

10

10

10

10

10

8

2+

2-

+

10 protons 10 neutrons and 9 electrons

10 protons 12 neutrons and 8 electrons

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 witha mass spectrometer

Figure 2-14

A mass spectrometer and a mass spectrum


2 5 atomic mass

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

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

Main Group

Lanthanides and Actinides

The Periodic table


2 7 the concept of the mole and the avogadro constant

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

  • 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)

NA = 6.02214179 x 1023 mol-1

  • Molar mass, M: the mass ofone mole of substance, from a table of atomic masses

  • Exp: the molar mass of lithium is 6.941 g/mol Li


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