Chapter Two Lecture

Atomic Theory of Matter

The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.

Dalton’s Postulates

Each element is composed of extremely small particles called atoms.

Still true!

Dalton’s Postulates

All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

True except for the discovery of isotopes!

Dalton’s Postulates

Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

Still true!

Dalton’s Postulates

Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

Still true!

Law of Constant CompositionJoseph Proust (1754–1826)

• Also known as the law of definite proportions.

• The elemental composition of a pure substance never varies.

Law of Conservation of Mass

The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

The Electron

• Streams of negatively charged particles were found to emanate from cathode tubes.

• J. J. Thompson is credited with their discovery (1897).

J.J. Thompson’s Cathode Tube the electron

The Electron

Thompson measured the charge/mass ratio of the electron to be 1.76 × 108 coulombs/g.

Millikan Oil Drop Experiment

Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.

Millikan Oil Drop Experiment

Robert Millikan (University of Chicago) determined the charge on the electron in 1909.

Radioactivity:

• The spontaneous emission of radiation by an atom.

• First observed by Henri Becquerel.• Also studied by Marie and Pierre Curie.

Radioactivity• Three types of radiation were discovered by

Ernest Rutherford:α particlesβ particles γ rays

The Atom, circa 1900:

• “Plum pudding” model, put forward by Thompson.

• Positive sphere of matter with negative electrons imbedded in it.

Discovery of the Nucleus

Ernest Rutherford shot α particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

The Nuclear Atom

Since some particles were deflected at large angles, Thompson’s model could not be correct.

The Nuclear Atom• Rutherford postulated a very small,

dense nucleus with the electrons around the outside of the atom.

• Most of the volume of the atom is empty space.

Rutherford’s Gold Foilthe nucleus

Other Subatomic Particles

• Protons were discovered by Rutherford in 1919.

• Neutrons were discovered by James Chadwick in 1932.

Subatomic Particles• Protons and electrons are the only particles that

have a charge.• Protons and neutrons have essentially the same

mass.• The mass of an electron is so small we ignore it.

Symbols of Elements

Elements are symbolized by one or two letters.

Atomic Number

All atoms of the same element have the same number of protons:

The atomic number (Z)

Atomic Mass

The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

Isotopes:

• Atoms of the same element with different masses.• Isotopes have different numbers of neutrons.

116C

126C

136C

146C

SAMPLE EXERCISE 2.2 Determining the Number of Subatomic Particles in Atoms

How many protons, neutrons, and electrons are in (a) an atom of 197Au (b) an atom of strontium-90?

SAMPLE EXERCISE 2.2 Determining the Number of Subatomic Particles in Atoms

How many protons, neutrons, and electrons are in (a) an atom of 197Au (b) an atom of strontium-90?

Solution (a) The superscript 197 is the mass number, the sum of the number of protons plus the number of neutrons. According to the list of elements given on the front inside cover, gold has an atomic number of 79. Consequently, an atom of 197Au has 79 protons, 79 electrons, and 197 – 79 = 118 neutrons. (b) The atomic number of strontium (listed on the front inside cover) is 38. Thus, all atoms of this element have 38 protons and 38 electrons. The strontium-90 isotope has 90 – 38 = 52 neutrons.

Atomic Mass

Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.

Average Mass

• Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations.

• Average mass is calculated from the isotopes of an element weighted by their relative abundances.

SAMPLE EXERCISE 2.4

Naturally occurring chlorine is 75.78% 35Cl, which has an atomic mass of 34.969 amu, and 24.22% 37Cl, which has an atomic mass of 36.966 amu. Calculate the average atomic mass (that is, the atomic weight) of chlorine.

SAMPLE EXERCISE 2.4

Naturally occurring chlorine is 75.78% 35Cl, which has an atomic mass of 34.969 amu, and 24.22% 37Cl, which has an atomic mass of 36.966 amu. Calculate the average atomic mass (that is, the atomic weight) of chlorine.

Solution The average atomic mass is found by multiplying the abundance of each isotope by its atomic mass and summing these products.

SAMPLE EXERCISE 2.4

Naturally occurring chlorine is 75.78% 35Cl, which has an atomic mass of 34.969 amu, and 24.22% 37Cl, which has an atomic mass of 36.966 amu. Calculate the average atomic mass (that is, the atomic weight) of chlorine.

Solution The average atomic mass is found by multiplying the abundance of each isotope by its atomic mass and summing these products.

SAMPLE EXERCISE 2.4

Naturally occurring chlorine is 75.78% 35Cl, which has an atomic mass of 34.969 amu, and 24.22% 37Cl, which has an atomic mass of 36.966 amu. Calculate the average atomic mass (that is, the atomic weight) of chlorine.

Solution The average atomic mass is found by multiplying the abundance of each isotope by its atomic mass and summing these products.

This answer makes sense: The average atomic mass of Cl is between the masses of the two isotopes and is closer to the value of 35Cl, which is the more abundant isotope.

BELLWORK

Which of the following diagrams is most likely to represent an ionic compound, and which a molecular one? Explain your choice.

Periodic Table:

• A systematic catalog of elements.

• Elements are arranged in order of atomic number.

Periodicity

When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

Periodic Table• The rows on the periodic

chart are periods.• Columns are groups.• Elements in the same

group have similar chemical properties.

Groups

These five groups are known by their names.

Periodic Table

Nonmetals are on the right side of the periodic table (with the exception of H).

Periodic Table

Metalloids border the stair-step line (with the exception of Al and Po).

Periodic Table

Metals are on the left side of the chart.

SAMPLE EXERCISE 2.5 Using the Periodic Table

Which two of the following elements would you expect to show the greatest similarity in chemical and physical properties: B, Ca, F, He, Mg, P?

PRACTICE EXERCISELocate Na (sodium) and Br (bromine) on the periodic table. Give the atomic number of each, and label each a metal, metalloid, or nonmetal.

Chemical Formulas

The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.

Molecular Compounds

Molecular compounds are composed of molecules and almost always contain only nonmetals.

Diatomic Molecules

These seven elements occur naturally as molecules containing two atoms.

Types of Formulas

• Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.

• Molecular formulas give the exact number of atoms of each element in a compound.

SAMPLE EXERCISE 2.6 Relating Empirical and Molecular Formulas

Write the empirical formulas for the following molecules: (a) glucose, a substance also known as either blood sugar or dextrose, whose molecular formula is C6H12O6; (b) nitrous oxide, a substance used as an anesthetic and commonly called laughing gas, whose molecular formula is N2O.

PRACTICE EXERCISE

Give the empirical formula for the substance called diborane, whose molecular formula is B2H6.

Types of Formulas

• Structural formulas show the order in which atoms are bonded.

• Perspective drawings also show the three-dimensional array of atoms in a compound.

Ions

• When atoms lose or gain electrons, they become ions.– Cations are positive and are formed by

elements on the left side of the periodic chart.– Anions are negative and are formed by

elements on the right side of the periodic chart.

SAMPLE EXERCISE 2.7 Writing Chemical Symbols for Ions

Give the chemical symbol, including mass number, for each of the following ions: (a) The ion with 22 protons, 26 neutrons, and 19 electrons; (b) the ion of sulfur that has 16 neutrons and 18 electrons.

PRACTICE EXERCISE

How many protons and electrons does the Se2– ion possess?

SAMPLE EXERCISE 2.8 Predicting the Charges of Ions

Predict the charge expected for the most stable ion of barium and for the most stable ion of oxygen.

PRACTICE EXERCISE

Predict the charge expected for the most stable ion of aluminum and for the most stable ion of fluorine.

Ionic Bonds

Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

SAMPLE EXERCISE 2.9 Identifying Ionic and Molecular Compounds

Which of the following compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4?

PRACTICE EXERCISE

Which of the following compounds are molecular: CBr4, FeS, P4O6, PbF2 ?

Writing Formulas

• Because compounds are electrically neutral, one can determine the formula of a compound this way:– The charge on the cation becomes the subscript

on the anion.– The charge on the anion becomes the subscript

on the cation.– If these subscripts are not in the lowest whole-

number ratio, divide them by the greatest common factor.

Common Cations

Common Anions

If you have equal masses of the following metals, which will have

the largest number of atoms? • Pb, 207.2 g/mol• Au, 197 g/mol• Ag, 107.9 g/mol• Cu, 63.54 g/mol• Fe, 55.85 g/mol

If you have equal masses of the following metals, which will have

the largest number of atoms? • Pb, 207.2 g/mol• Au, 197 g/mol• Ag, 107.9 g/mol• Cu, 63.54 g/mol• Fe, 55.85 g/mol

The mass spectrum of chlorine is shown below. How many different molar masses are possible for Cl2?

• 1• 2• 3• 4• 5

The mass spectrum of chlorine is shown below. How many different molar masses are possible for Cl2?

• 1• 2• 3• 4• 5

Which molar mass will be most common among Cl2 molecules?

• 70 g/mol• 71 g/mol• 72 g/mol• 73 g/mol• 74 g/mol

Which molar mass will be most common among Cl2 molecules?

• 70 g/mol• 71 g/mol• 72 g/mol• 73 g/mol• 74 g/mol

Which combination is likely to produce an ionic compound?

• C and H• S and Cl• Ca and F• Br and I• Xe and F

Which combination is likely to produce an ionic compound?

• C and H• S and Cl• Ca and F• Br and I• Xe and F

What are the ions present in the compound

(NH4)2SO4 ? • NH3, H2, and SO2

• N3-, H+, S2-, O2-

• NH42+ and SO4

-

• NH4+ and SO4

2-

• (NH4+)2 and SO4

2-

What are the ions present in the compound

(NH4)2SO4 ? • NH3, H2, and SO2

• N3-, H+, S2-, O2-

• NH42+ and SO4

-

• NH4+ and SO4

2-

• (NH4+)2 and SO4

2-

Bellwork Write the molecular and structural formulas for the compounds represented

Common Cations

Common Anions

Inorganic Nomenclature

• Write the name of the cation.• If the anion is an element, change its ending

to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion.

• If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

SAMPLE EXERCISE 2.12 Determining the Names of Ionic Compounds from Their Formulas

Name the following compounds: (a) K2SO4 , (b) Ba(OH)2 , (c) FeCl3.

(a) potassium sulfate. (b) barium hydroxide. (c) iron(III) chloride or ferric chloride.

PRACTICE EXERCISE

Name the following compounds: (a) NH4Br, (b) Cr2O3, (c) Ca(NO3)2.

(a) ammonium bromide, (b) chromium(III) oxide, (c) calcium nitrate

SAMPLE EXERCISE 2.13 Determining the Formulas of Ionic Compounds from

Their Names Write the chemical formulas for the following compounds: (a) potassium sulfide, (b) calcium hydrogen carbonate, (c) nickel(II) perchlorate.

(a) K2S. (b) Ca(HCO3)2. (c) Ni(ClO4)2.

PRACTICE EXERCISE

Give the chemical formula for (a) magnesium sulfate, (b) silver sulfide, (c) lead(II) nitrate.

(a) MgSO4 , (b) Ag2S, (c) Pb(NO3)2

Patterns in Oxyanion Nomenclature

• When there are two oxyanions involving the same element:– The one with fewer oxygens ends in -ite

• NO2− : nitrite; SO3

2− : sulfite

– The one with more oxygens ends in -ate• NO3

− : nitrate; SO42− : sulfate

Patterns in Oxyanion • The one with the second fewest oxygens ends in -ite

ClO2− : chlorite

• The one with the second most oxygens ends in -ateClO3

− : chlorate

Patterns in Oxyanion Nomenclature• The one with the fewest oxygens has the prefix hypo- and

ends in -iteClO− : hypochlorite

• The one with the most oxygens has the prefix per- and ends in -ate

ClO4− : perchlorate

SAMPLE EXERCISE 2.11 Determining the Formula of an Oxyanion from Its Name

Based on the formula for the sulfate ion, predict the formula for (a) the selenate ion and (b) the selenite ion. (Sulfur and selenium are both members of group 6A and form analogous oxyanions.)

PRACTICE EXERCISE

The formula for the bromate ion is analogous to that for the chlorate ion. Write the formula for the hypobromite and perbromate ions.

Acid Nomenclature• If the anion in the acid

ends in -ide, change the ending to -ic acid and add the prefix hydro- :– HCl: hydrochloric acid– HBr: hydrobromic acid– HI: hydroiodic acid

Acid Nomenclature• If the anion in the acid

ends in -ate, change the ending to -ic acid:– HClO3: chloric acid

– HClO4: perchloric acid

Acid Nomenclature• If the anion in the acid

ends in -ite, change the ending to -ous acid:– HClO: hypochlorous

acid– HClO2: chlorous acid

SAMPLE EXERCISE 2.14 Relating the Names and Formulas of Acids

Name the following acids: (a) HCN, (b) HNO3, (c) H2SO4, (d) H2SO3.

(a)hydrocyanic acid (b)nitric acid (c)sulfuric acid. (d)sulfurous acid

PRACTICE EXERCISE

Give the chemical formulas for (a) hydrobromic acid, (b) carbonic acid.

(a) HBr, (b) H2CO3

Nomenclature of Binary Compounds

• The less electronegative atom is usually listed first.

• A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however.)

Nomenclature of Binary Compounds

• The ending on the more electronegative element is changed to -ide.

– CO2: carbon dioxide– CCl4: carbon tetrachloride

Nomenclature of Binary Compounds

If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one:

N2O5: dinitrogen pentoxide

SAMPLE EXERCISE 2.15 Relating the Names and Formulas of

Binary Molecular Compounds

Name the following compounds: (a) SO2, (b) PCl5, (c) N2O3.

(a) sulfur dioxide (b) phosphorus pentachloride (c) dinitrogen trioxide.

PRACTICE EXERCISE

Give the chemical formula for (a) silicon tetrabromide, (b) disulfur dichloride.

(a) SiBr4, (b) S2Cl2

(a) Alkanes contain only carbon and hydrogen, and each carbon atom is attached to four other atoms.

Consider the alkane called pentane. (a) Assuming that the carbon atoms are in a straight line, write a structural formula for pentane. (b) What is the molecular formula for pentane?

PRACTICE EXERCISEButane is the alkane with four carbon atoms. (a) What is the molecular formula of butane? (b) What are the name and molecular formula of an alcohol derived from butane?

(a) C4H10 (b) butanol, C4H10O or C4H9OH

(b) C5H12