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Chapter 2 Atoms, Molecules, and Ions Chemistry, The Central Science, 13th edition Theodore L. Brown, H. Eugene LeMay, Jr., Bruce E. Bursten, Murphy, Woodward,

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Presentation on theme: "Chapter 2 Atoms, Molecules, and Ions Chemistry, The Central Science, 13th edition Theodore L. Brown, H. Eugene LeMay, Jr., Bruce E. Bursten, Murphy, Woodward,"— Presentation transcript:

1 Chapter 2 Atoms, Molecules, and Ions Chemistry, The Central Science, 13th edition Theodore L. Brown, H. Eugene LeMay, Jr., Bruce E. Bursten, Murphy, Woodward, and Stoltzfus.

2 2.1 The Atomic Theory of Matter

3 Greek Philosophers  Early philosophers debated the fundamental “stuff” of which the universe was made  Democritus’ idea of the world was based on “atomos” meaning indivisible or uncuttable  Later Plato and Aristotle formulated that there can be no uncuttable “things” and they successfully propagated their theory and therefor became the dominant theory for centuries

4 John Dalton  The atom reemerged in the 17 th century  The ground work was laid by John Dalton when he noticed that elements that reacted with other elements to form new compounds  Dalton’s atomic theory was based on 4 postulates

5 Dalton’s Atomic Theory  Explains several laws that were known during Dalton’s time, including the law of constant composition.  In a given compound, the relative numbers and kinds of atoms are constant  The Law of Conservation of Mass  The total mass of materials present after a chemical reaction is the same as the total mass present before the reaction  The Law of Multiple Proportions  If two elements A and B combine to form more than one compound, the masses of B that can combine with given mass of A are in the ratio of small whole numbers

6 Dalton’s Atomic Theory

7 Dalton’s Theory Correct?

8 Give it some thought  Compound A contains 1.333 g of oxygen per gram of carbon, whereas compound B contains 2.666 g of oxygen per gram of carbon.  What chemical law does this data illustrate?  Law of Multiple Proportions  If compound A has an equal number of oxygen and carbon atoms, what can we conclude about the composition of compound B?  There are twice as many oxygen atoms in compound B than there is in compound A

9 2.2 The Discovery of the Atomic Structure

10 Subatomic Particles  Dalton based his theory on observations made in the laboratory  He and those who followed had no direct evidence of the atom  As time progressed, scientists began to probe the nature of matter and started to discover subatomic particles  We will see that the atom is composed in part by electrically charged particles  Keep in mind as we continue that same charges repel one another, whereas particles with unlike charges attract one another

11 Cathode Rays and Electrons (e-)  During the 1800’s scientists experimented (including Thomson) with evacuated glass tubes with electrodes inserted at both ends  Once a charge was applied a radiation between the electrodes was produced  Cathode rays emanated from the negative end and traveled to the positive end  The rays were unseen but caused certain materials to fluoresce  The rays were tested with magnets and electrically charged rods and it was found that the cathode rays were negatively charged  The identity of the new particle was the same regardless of the gas used  This new “thing” was called the electron  Thomson tested the beam and found that there was 1.76 x 10 8 coulombs per gram



14 Give it some thought  Thomson observed that the cathode rays produced in the cathode–ray tube behaved identically, regardless of the particular metal used as cathode. What is significance of this observation?  All atoms have the same subatomic particles called electrons!

15 Millikan Oil Drop Experiment  Once the charge-to-mass ratio of the electron was known, Milikan was then able to experimentally figure out the mass of the electron  Using an experiment similar to the one pictured to the right he calculated the mass of the electron to be 9.10 x 10 -28 g by solving the charge of a single electron  This showed that the electron has a mass of about 2000 times less than hydrogen!

16 Radioactivity  In 1896, Henri Becquerel discovered that uranium emitted radiation.  Spontaneous emission is known as radioactivity  He concluded that the source or the radiation was the uranium atoms  Studies done by Ernest Rutherford showed that there were 3 types of radiation  Alpha (α)  Beta (β)  Gamma (γ)  Alpha and beta radiation was shown to be bent by an electric field (but in different directions) while gamma radiation was unaffected by it

17 α and β Rays  Both considered fast moving particles  Beta rays were shown to be the radioactive equivalent of cathode rays  Attracted to positively charged plates!  Charge of -1  The alpha particles were shown to have a positive charge  Charge of +2  Has a mass 7400 times that of an electron  Gamma radiation is high-energy radiation similar to X-rays and does not consist of particles and carries no charge

18 3 Types of Radiation

19 Thomson Model  Thomson in the early 1900’s reasoned that since the electrons are such a small portion of the mass of the atom that it must also only make up a small portion of the atom  Thomson proposed that that the electrons were embedded in the atom like raisons in pudding  The “pudding” having a positive charge

20 Rutherford Atomic Model  Rutherford was studying the angles at which alpha particles were being deflected  He discovered that most particles passed straight through the gold foil while only a small amount of the particles were deflected or even bounced back  Rutherford explained the result by postulating the nuclear model of the atom, which most of the mass of each gold atom resided in the nucleus  He postulated that the majority of the space in an atom was empty and that the charge of the nucleus was positive

21 Protons and Neutrons  Subsequent experiments led to the discovery of positive particles (protons) and neutral particles (neutrons) in the nucleus.  Protons were discovered in 1919 by Rutherford and neutrons in 1932 by British scientist James Chadwick (1891–1972). Thus, the atom is composed of electrons, protons, and neutrons.

22 Give it Some Thought  What happens to most of the alpha particles that strike the gold foil in Rutherford’s experiment?  They passed straight through  Why do they behave that way?  The atom is mostly empty space

23 2.3 The Modern View of Atomic Structure

24 The Modern View of Atomic Structure  The charge of an electron is -1.602 * 10 -19 C.  The charge of a proton is opposite in sign but equal in magnitude to that of an electron: +1.602 * 10 -19 C.  The quantity 1.602 * 10 -19 C is called the electronic charge.  For convenience, the charges of atomic and subatomic particles are usually expressed as multiples of this charge rather than in coulombs.  The charge of an electron is 1- and that of a proton is 1+.  Neutrons are electrically neutral  Every atom has an equal number of electrons and protons, so atoms have no net electrical charge.

25 Atomic Diameter  Protons and neutrons reside in the center of the atom and take up only a small volume of the atom  Most atoms have diameters between 1 * 10-10 m 1100 pm2 and 5 * 10-10 m 1500 pm2. A convenient non–SI unit of length used for atomic dimensions is the angstrom 1A° 2, where 1 A° = 1 * 10-10 m. Thus, atoms have diameters of approximately 1 - 5 A°. The diameter of a chlorine atom, for example, is 200 pm, or 2.0 A°.

26 Give it Some Thought  If an atom has 15 protons, how many electrons does it have?  Phosphorous  Where do the protons reside in an atom?  The nucleus

27 Subatomic Particles

28 Sample Exercise 2.1  The diameter of a U.S. dime is 17.9 mm, and the diameter of a silver atom is 2.88 A°. How many silver atoms could be arranged side by side across the diameter of a dime?


30 Atomic Numbers, Mass Numbers, and Isotopes  The atoms of each element have a characteristic number of protons. The number of protons in an atom of any particular element is called that element’s atomic number.  Because an atom has no net electrical charge, the number of electrons it contains must equal the number of protons.  The atomic number is indicated by the subscript; the superscript, called, the mass number is the number of protons and neutrons

31 Isotopes  Isotopes differ by the number of neutrons but contain the same number of protons



34 2.4 Atomic Weights

35 Atomics Mass Scale  Scientists in the 19 th century were aware that atoms of different elements have different masses  They knew that in water that there was 88.9 g of oxygen for every 11.1 grams of hydrogen in 100 grams of water  Today we can determine the masses of individual atoms with a high degree of accuracy.  The atomic mass unit is a convenient way when dealing with very small masses  i.e 1 mole of carbon twelve weights 12 amu

36 Atomic Weight  Atomic weight is the natural mixture of isotopes found in nature  To calculate the atomic weight of an element is to use the following equation:  Atomic Weight = Σ[(isotope mass) x (fractional isotope abundance)]  Σ sigma means over all masses

37 Give it Some Thought  A particular atom of chromium has a mass of 52.94 amu, whereas the atomic weight of chromium is given as 51.99 amu. Explain the difference in the two masses.


39 Mass Spectrometers

40 Some Pictures From Mr. Hunter’s Graduate Research

41 Actual Data Gathered by Mr. Hunter Why are there bands over a certain mass range and not just a line? The mass spectrometer is an energy filter that allows only stable trajectories of ions through. Therefore, all ions of the same mass have a range of energies (different velocity) coming from the ion source.

42 Book Work in Class  Pages 74-75  Answer only red questions for section 2.1-2.4

43 2.5 The Periodic Table

44 The Periodic Table  As the list of elements began to grow throughout history, certain elements were noticed to have similar traits as other elements.  For example, Na and K had similar properties while He and Ne also had similar properties to one another.  These similarities became known as periodic trends or families

45 Periodic Table Periods= horizontal rows Families (groups) = vertical columns

46 Common Families

47 Metals vs Nonmetals Metals are found to the left of the stair-step line Non-metals are found to the right of the stair-step line Metalloids are found along the stair-step line


49 2.6 Molecules and Molecular Compounds

50 Molecules and Molecular Compounds All of the pictured molecules are in their chemical formulas except for one Oxygen and hydrogen exist as diatomic elements, i.e molecule Molecules that contain more than one type of atom are called molecular compounds

51 Molecular and Empirical Formulas  Chemical formulas that indicated the exact number of atoms in a molecule are called molecular formulas  Chemical formulas that indicate the relative number of atoms in a molecule are called empirical formulas


53 Picturing Molecules  Structural formulas shows which atoms are attached to which  Atoms can be represented by their symbols and bonds by lines  This type of modeling does not show the actual geometry of the molecule  The Ball-and Stick model has the advantage of showing the actual geometry of the atoms in relation to one another  The space filling model shows the relative sizes of atoms and their geometries


55 Book Work  Page 75-76  Answer only red questions for section 2.5-2.6

56 2.7 Ions and Ionic Compounds

57 Ions and Ionic Compounds  An ion is a charged particle  A cation is a positive particle (one or more less electrons)  An anion is a negative particle (one or more extra electrons)  Ions form because certain elements have a weak attraction to the outer electrons (valence electrons)  Other elements have a strong attraction for their electrons and surrounding elements electrons  The stronger pull can take the weaker element’s electron  This transfer of electrons causes ions to form  In addition to simple ions there are polyatomic ions such as ammonium and sulfate  These ions are made of at least 2 elements


59 Predicting Ionic Charges  Elements like to have a stable number of electrons.  The most stable elements on the periodic table are the noble gases because they have a full octet of electrons  Elements gain or lose electrons to get to this stable state  Nonmetals tend to gain electrons  Metals tend to lose electrons  Half Filled shells are more stable that partially filled shells  Meaning, electrons can move from one shell to another to stabilize the ion  This also means that s-orbital electrons can be taken instead of f- or d-shell electrons

60 Practice Problem  Predict the charge states of:  Fe  Sn  Sc  Ti  V




64 Ionic Compounds Ionic compounds are a compounds made up of anions and cations Ionic compounds are generally metals and nonmetals.



67 Criss-Cross Method  Since there are not discrete ionic “molecules” we are only able to write empirical formulas for these compounds  The criss-cross method take the charges of each ion and puts it down at the subscript of the opposite ion


69 2.8 Naming Inorganic Compounds

70 Naming Inorganic Compounds  Chemical Nomenclature comes from the latin word nomen (name) and calare (to call)  The rules of chemical nomenclature are based on the division of substances into categories  Organic compounds contain carbon and hydrogen often with oxygen, nitrogen, and other elements  All other compounds are inorganic compounds

71 Names and Formulas of Ionic Compounds Most metals that form cations can form different oxidation states. These are usually transition metals







78 Oxyanions- an anion containing one or more atoms bonded to another element (as in the sulfate and carbonate ions) Oxyacids- an acid “see above“



81 Names and Formulas of Acids


83 Names and Formulas of Binary Molecular Compounds  The name of the element farther to the left in the periodic table (closest to the metals) is usually written first. An exception occurs when the compound contains oxygen and chlorine, bromine, or iodine (any halogen except fluorine), in which case oxygen is written last.  If both elements are in the same group, the one closer to the bottom of the table is  named first.  The name of the second element is given an -ide ending.  Greek prefixes indicate the number of atoms of each element.  (Exception: The prefix mono- is never used with the first element.) When the prefix ends in a or o and the name of the second element begins with a vowel, the a or o of the prefix is often dropped.

84 Give it Some thought  Is SOCl 2 a binary compound?


86 Book Problems  Answer on red question for section 2.7-2.8

87 2.9 Some Simple Organic Compounds

88 Organic Chemistry  The study of compounds containing carbon is called organic chemistry  Organic chemistry is a very important branch of chemistry encompassing a large portion or what chemists do  Being able to name these compounds is very important for the standardization of chemistry experiments  Hydrocarbons are composed of only carbon and hydrogen  Alkanes are the simplest of hydrocarbons  Alkanes always end in -ane

89 Some Derivatives of Alkanes  Functional groups can be added to organic molecules to change their chemical properties  Alcohols are made by adding an –OH group in place of a hydrogen  An alcohol has the ending –ol  Numbers are used to determine what carbon the functional group is attached  Always use numbers that make it the smallest possible number or smallest possible addend

90 Isomers  Isomers are compounds that have the same molecular formulas but a different chemical arrangement  The most common type of isomer are structural isomers

91 Other Functional Groups  There are hundreds of different functional groups in organic chemistry  Common functional groups are:  Halogens  Hydroxyl (-OH)  Carboxylic acid (-COOH)  Oxygen insertions  Carbonyl groups (-CO)  Alkane groups

92 Common Functional Groups

93 Give it some thought  Draw the structural formulas of two isomers of butane, C 4 H 10

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