TOPIC II: ELEMENTS AND COMPOUNDS

• Nuclear Atom

• Isotopes

• Periodic Table of Elements

• Molecules and Compounds

• Naming Simple Ionic Compounds

• Naming Binary Molecules

Kotz & Treichel Ch 2,3

Chapter 2: Atoms and Elements

1. History of discovery of the atomic nature of all matter and the various particles found within the atom: Kotz, 2.1-2.2, excellent reading on all topics. Saunders CD-ROM: animated film clips

Reading, viewing assignment

Lecture Topics, Kotz, 2.3-2.8

Atomic Structure, Atomic Number, Mass

All matter (anything that has mass and occupiesvolume) can be classified as:

• an element (basic building blocks of nature: H, O, Au )

• a compound (made up of two or more elements)

• a mixture (any physical combination of the above)

Elements, the simplest forms of matter, are composed of unique tiny particles called atoms.

Atomic Structure

The atom itself is composed of three types of “subatomic particles”, the proton (p), the neutron (n),and the electron (e).

Each element has its own unique pairings of thesethree particles.

It is the number and the placement of these particles which gives rise to the different properties exhibited by each element.

proton neutron electron

mass, g 1.673 E-24 g 1.675 E-24 g 9.11 E-28 g

mass, amu 1.007 amu 1.009 amu 0.000549 amu

comparative mass 1 1 0

relative charge 1 0 -1

Comparative Mass, Charge: Nuclear Particles

Nuclear Particle Location Within the Atom

1. Protons and Neutrons: “Nucleus of Atom” compact positive mass in center of atom

“all” of the mass, negligible volume (10-2 pm)

2. Electrons: “Outside the Nucleus” cloud of negative charge

“all” of the volume (103 pm), negligible mass

THE “NUCLEAR” ATOM

Nucleus, C atom

Electron cloud

.

ATOMIC SYMBOLS

Each atom of an element can be represented by a symbol that describes how many protons, neutronsand electrons are contained in this basic unit:

XA

Z

Mass number, A:total #, p + n

Atomic Number, Z: #p (equals #e)

Elementalsymbol

Accordingly, A, the mass number of the atom, representsboth the total number of nuclear particles (p + n) and the approximate mass of the atom (in amu’s)

Because atoms (and the subatomic particles) are so tiny,a relative mass scale was setup to describe atomicweights in a convenient numerical range.

The atom of Carbon which contains 6 protons and 6 neutrons in its nucleus is assigned the weight of 12.00 amu (atomic weight units), which essentiallymakes the mass of each proton and neutron 1.00 amu.

The Mass Number, A

Z, The Atomic Number

All known elements are listed in the familiar “PeriodicTable of the Elements” in order of increasing atomicnumber, found usually in the upper right hand cornerabove each element’s symbol.

The atomic number represents the number of protonsin the nucleus of every atom of that element.

Since every atom is electrically neutral, the number ofprotons in the nucleus represents the total positivecharge of the nucleus, which is exactly balanced by the total number of negatively charged electrons outside the nucleus.

Examples, Atomic symbols

F U19

9

238

92

X

A

Z

Atomic Mass #, p + n

Atomic #, p (or e)

9p, 9e

19 - 9p = 10n 238 - 92p = 146n

92p, 92e

Element

Group Work 2.1: Atomic Symbols

Br35

80

Pb82

207

Cr52

24

??40

20

p's n's e's

Isotopes of the Elements

For a given element, the number of protons and electrons is a fixed value and determines which element is being described.

The number of neutrons in the atom of a givenelement is not fixed, and several different atoms ofa given element are generally found, varying in atomic mass due to differing numbers of neutrons.

The different atoms of a given element which vary byneutron count and by mass are described as “isotopes”of that element.

Isotopic Symbols

Isotopic symbols represent specific isotopic forms of an element. Consider hydrogen:

H H H1 1 1

1 2 3

1 p0 n

1 p1 n

1 p2 n

"protium" "deuterium" "tritium"

Atomic Mass, revisited:

The atomic mass of any isotope of an element can beapproximated by a simple addition of the number ofp’s and n’s in the nucleus.

However, naturally occurring samples of any element generally include several different isotopes of different atomic masses.

The atomic mass value given for each element (asfound in your Periodic Table) is a weighted average of all the isotopes.

The given atomic mass for any element will beclosest to the most abundant element in most cases. Consider below the calculation of theatomic mass of magnesium, 24.305 amu:

Average Mass Element, amu =

(mass, isotope A X % abundance A)

+ (mass, isotope B X % abundance B)

+ (mass, isotope C X % abundance C)......

General Method of Calculation:

Calculation of average atomic mass of Magnesium:

Mg Mg12 12 12

24 25 26Mg

23.9850 amu

78.99%

24.9858 amu

10.00 %

25.9826 amu

11.01%

(23.9850 x .7899) + (24.9858 x .1000) + (25.9826 x .1101)

= (18.95) + (2.499) + (2.861) = 24.31 amu

You may detect a discrepancy between the mass, in amu’s, of individual protons and neutrons (as given in the first table presented, lecture 3) and the given masses of the Mg isotopes:

Note:

Mass, proton: 1.007 amuMass, neutron, 1.009 amu

Mass, 24 Mg , 23.9850 amu 12

1.007 amu X 12 = 12.0841.009 amu X 12= 12.108 24.192

Why Don’t the Numbers Add Up?

The mass of the individual particles, p and n,were determined by methods which measured individual particles not bound up in the nucleus.

When these particles are packed into the nucleusof an atom, a mass loss (or “mass defect”) alwaysoccurs. This loss is thought to be the result ofa matter to energy conversion which holds togetherthe nucleus, called the “binding energy”...

(See Chapter 24, p 1099)

In the nucleus of atoms, p’s and n’s show a mass very close to 1.00 amu rather than 1.007 or 1.009.

The Periodic Table

• Lists all known elements in order of increasingatomic number, left to right and top to bottom

•Arranged so that elements of similar chemicalproperties fall into the same column (called a groupor family)

•Each horizontal line(called a period) represents elements of a complete range of chemical properties.

• The beginning of each “period” features a very active metal

•The middle of the period includes increasingly less active metals

•The right hand side of the period includes elementsbecoming increasingly non-metallic

•The end of the period features an inert, unreactiveelement found only in the gas state

Each line of the table features a complete swing from reactive metal to non metal to inert gas.

The action is repeated in each successive period, named after the action of a pendulum, in which each revolution or swing is also called a period or “repeated occurrence , from beginning to end and then back again”

1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A1 H He2 Li Be B C N O F Ne3 Na Mg Al Si P S Cl Ar4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe6 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn7 Fr Ra Ac Rf Db Sg Bh Hs Mt

metals

metalloids

Non metals

Noble gases

1A 2A 3A 4A 5A 6A 7A 8AH HeLi Be B C N O F NeNa Mg Al Si P S Cl ArK Ca Ga Ge As Se Br KrRb Sr In Sn Sb Te I XeCs Ba Tl Pb Bi Po At RnFr Ra

Main Group or “Representative Elements”

Alkali metals 1A (except H)

Alkaline earthmetals 2A

Halogens 7A

Noble gasesinert gasesrare gases 8A (0)

3B 4B 5B 6B 7B 8B 8B 8B 1B 2B1234 Sc Ti V Cr Mn Fe Co Ni Cu Zn5 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd6 La Hf Ta W Re Os Ir Pt Au Hg7 Ac Rf Db Sg Bh Hs Mt

Transition Metals3 columns 8B:more alike “across”than “down”

Non- naturally occurring, newly found in lab

La #57 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf #72Ac #89 Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf #104

The inner Transition Metals: Found beneath the main body of the PT, but belonging to Period 6 and 7:

#58

#90 #103

#71“Lanthanides”

“Actinides”

1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A1 H2 He

2 Li Be B C N2 O2 F2 Ne

3 Na Mg Al Si P S Cl2 Ar

4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br2 Kr

5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I2 Xe6 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn7 Fr Ra Ac Rf Db Sg Bh Hs Mt

solids liquids gases

Note the seven common elements found in nature as “diatomic molecules”

Molecules and Compounds, Chapter 3

Molecules are made by the bonding of two or more atoms together into a independent, uncharged unit.The simplest molecules are the “diatomic elements”:

H2 N2 O2 F2 Cl2 Br2 I2

Atoms of the elements come together to form molecules and compounds.

Lecture Topics, Kotz, 3.1-3.5

Compounds of the elements may be described as “molecular” or “ionic” in nature, depending on the nature of the elements involved:

Molecular compounds are typically formed between non-metals, bonded together in a single “molecular”unit: CO2 H2O NH3 PBr3 CH4 SO3

Ionic compounds are formed when metals bond to nonmetals. The metal and the nonmetal components exist as individual charged units called “ions”:

NaCl (Na+ Cl-) Fe(OH)3 [Fe3+, 3 (OH)- ]

CuSO4 (Cu2+ SO42- )

Formulas for Molecules and

Compounds All compounds and molecules are represented by a formula which summarizes number of atoms of each element present: H2SO4 CH3CH2OH NaCl

C12H22O11 F2 Fe2(SO4)3 (NH4)2CO3 H2O

IONIC: formula represents the simplest ratio of positive to negative ions found in a sample of the compound...

MOLECULAR: formula represents all atoms in formula bonded together in a single unit called a molecule...

TYPES OF FORMULAS

• CONDENSED FORMULAS: summation of all atoms in formula, alcohol: C2H6O

• STRUCTURAL FORMULAS: clue to connectivity in compound, alcohol: CH3CH2OH

H

O

H

C C

H

H

H H

H

O

H S

O

O

OO

HH

Cl -Na +

Na + Cl -

Na +

Cl -

Cl -

Na +

NAMING MOLECULES AND COMPOUNDS

There are many rules for naming many types of compounds: To name any compound, you must first recognize its “type.”

“Guidelines”:

If the compound formula starts with a nonmetal or metalloid,it is a molecular type compound

If the formula starts with a metal, consider it an ionic compound, If the formula starts with H, it is an acid

We’ll name in this lesson two specific types:

•“binary molecules” (two non-metals in formula ) CO2 PCl3 SO3 AsF3

•“ionic salts and bases” ( cation / anion combinations)

cation: some metal ion or NH4+

anion: OH- or O2- (a “base”) all other anions: (“salts”)

K2SO4 NH4Cl NaOH MgO FeBr3

Acids (H written first) will be done later...

GROUP WORK 2.2: Type of Compound Acid, Base, Salt, or Molecular?

Compound Type Compound Type

Pb(NO3)2 HNO3

HClO4 Cl2O7

Pb(OH)4 CrSO4

SbH3 SiCl4

CaSiO4 PF5

Fe2O3 Fe3(PO4)2

Before naming ionic compounds, let’s review cations and anions, and examine charges and formulas we need to know.....

POSITIVE IONS: Cations

CATIONS: positively charged ions; monoatomic cations are formed from metals which have LOST one or more electrons in compound formation:

- 1 eNa (11p, 11e) --------> Na+ (11p, 10e) (all Group 1A)

-2eCa (20p, 20e) --------> Ca2+ (20p, 18e) (all Group 2A)

1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A12 Li Be3 Na Mg Al4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga5 Rb Sr Y Zr Nb Mo Ru Rh Pd Ag Cd In Sn6 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi

Common Metals, Fixed vs Variable Charge

Metals which form single cation

Metals which form several cations

Naming Cations: Fixed Charge Metals

When the metal in the salt or base exhibits only one charge and forms only one cation,the name of the cation is identical with that of the metal:

Na+ sodium cation Mg2+ magnesium cationAl3+ aluminum cation Ag+ silver cation

“One Charge Only”

1+ 2+

1+

3+2+ 3+

1A 2A 3B 1B 2B 3A1

2 Li+ Be2+

3 Na+ Mg2+ (3+) Al3+

4 K+ Ca2+ (3+) Zn2+

5 Rb+ Sr2+ (3+) Ag+ Cd2+

6 Cs+ Ba2+ (3+)

Metals Forming Several Cations

All other common metals form cations resulting from the loss of a variable number of electrons (depending on the circumstances of the reaction).

An examination of electronic structure (Unit 3) willjustify all charges, single or multiple; now we simply must recognize “which is which”.........

Naming Cations: Metals with Variable Charges

When a metal is known to form several differentcations of different charges, then the name of the cation must include a Roman Numeral indicatingthe charge of the ion:

Fe2+ Iron(II) cation Cu+ Copper(I) cationFe3+ Iron(III) cation Cu2+ Copper(II) cation

Sn2+ Tin(II) cation Bi3+ Bismuth(III) cation

Typically Encountered Cations, Variable Charge Metals

Cr2,3+

Fe2,3+

Co2,3+

Ni, Mn 2+

Cu1,2+

Hg2+, Hg22+

Sn, Pb 2+

Bi3+

Maximum (if not always common) charge on all metals isgiven by group number...

NEGATIVE IONS: Anions

Monoatomic ANIONS:

Single nonmetallic atoms which have gained one or more electrons in a chemical reaction and become negatively charged ions : +3eN (7p, 7e) --------> N3- (7p, 10e) (all Group 5A) +2eO(8p, 8e) --------> O2- (8p, 10e) (all Group 6A) +1eF (9p, 9e) --------> F1- (9p, 10e) (all Group 7A)

1A 5A 6A 7A

1 H-

2 N3- O2- F-

3 P3- S2- Cl-

4 Se2- Br-

5 I-

6

3-

Monoatomic Anions: Name, Charge

2- 1-

Hydride Nitride Oxide Phosphide Sulfide Selenide

FluorideChlorideBromideIodide

“ide”“ide”

FORMING IONIC COMPOUNDS

Na+ Cl- NaCl

Na+ S2- Na2S

Na+ P3- Na3P

Ba2+ Cl- BaCl2

Ba2+ S2- BaS

Ba2+ P3- Ba3P2

Make sure charges balance; cross multiply whencation and anion charges are different:

GROUP WORK 2.3:Form Compound, Name

• CATION, ANION

• Mg2+ H -

• Fe3+ S2-

• Al3+ P3-

• Cd2+ F -

• Mn2+ I -

• FORMULA, NAME

POLYATOMIC IONS

CATIONS: only one common, ammonium ion, NH4+

ANIONS: negatively charged ions containing two or more elements; the knowledge of the formula and charge of the most common are basic to naming compounds and writing formulas.

One of the elements usually involved is oxygen; theion names end in “ate” or “ite” as follows...

Key Polyatomic Anion Formers:Know these!

Br, I same as Cl

4A 5A 6A 7A12 C N O F3 P S Cl45

Permanganate (7B)like Perchlorate

Chromate (6B) like Sulfate

Polyatomic Anions of C, 4A

Most common: CO32- carbonate

HCO31- hydrogen carbonate,

“bicarbonate”

Others: CH3CO2- acetate (“C2H3O2

-”) CN- cyanide

Polyatomic Anions of N, P 5A

Nitrogen: NO3- nitrate

NO2- nitrite

(Remember also: NH4+, ammonium; N3-, nitride)

Phosphorus: PO43- phosphate

HPO42- hydrogen phosphate

H2PO41- dihydrogen phosphate

(Remember also: P3-, phosphide)

GROUP WORK 2.4

• CATION, ANION

• Cr3+ CO32-

• Ni2+ CN-

• Zn2+ NO2-

• Bi3+ H2PO4-

• Pb2+ N3-

• FORMULA, NAME

Polyatomic Anions of O, S (6A) Cr (6B)

Oxygen: OH- hydroxide

Remember also: O2- oxide

Sulfur: SO42- sulfate

SO32- sulfite

HSO4 - hydrogen sulfate

HSO3 - hydrogen sulfite

Remember also: S2-, sulfide Chromium: CrO4

2- chromate

The “hydroxides” and “oxides” of the metallic elements are referred to as “bases”; all other ionic combinations are referred to as “salts”

“BASES” and “SALTS”

Bases:

Mg(OH)2

NaOH

CaO

Fe(OH)3

Salts:

MgCl2 MgHSO4 MgCO3

Na3PO4 NaNO2 Na2SO3

Ca(NO3)2 Ca3N2 CaSO4

Fe(CN)2 Fe(CH3CO2)3 Fe(H2PO4)2

Polyatomic Anions of Cl, Br, I (7A) Mn (7B)

Fluorine, F forms only the monatomic anion F-; Bromine, Br and Iodine, I form the same ions as chlorine, Cl:

Chlorine: ClO- hypochlorite ClO2

- chlorite ClO3

- chlorate ClO4

- perchlorate

Remember also: Cl-, Chloride

Manganese: MnO4- permanganate

SUMMARY, NAMING IONIC SALTS AND BASES

State name of the cation, then name of the anion.

Cations with a variable charge are named by adding a Roman numeral

Monoatomic anions are named by changing

theirelemental name to end in “ide”

Polyatomic anions (memorized, Table 3.1, p. 110) end in “ite” or “ate”...

GROUP WORK 2.5

FORMULA

NH4ClO

Cd(BrO2)2

Co(IO3)3

Ba(ClO4)2

KMnO4

Ag2CrO4

NAME

Naming Binary Molecular Compounds

All compounds beginning with a metal or ammonium are named as “ ionic compounds.”

Compounds containing only two elements (“binary”) in which both elements in the formula are a non-metal or metalloid are named in a different manner...

The change in nomenclature reflects the fact that these compounds are “molecular” and not “ionic” in nature!

• Name the first element in the formula

• Name the second element in the formula to end in “ide”:

carbide, nitride, phosphide, oxide, sulfide, fluoride, bromide, chloride, iodide •Add numerical prefixes to indicate more than one atom of the element in the formula:

di (2), tri (3), tetra (4), penta (5), hexa (6), hepta (7), octa(8)

Binary Molecular Nomenclature Method:

Typical Nomenclature

• NO2

• SF6

• ICl5

• N2O5

• CBr4

• SO3

• P2O3

• nitrogen dioxide• sulfur hexafluoride• iodine pentachloride• dinitrogen pentoxide• carbon tetrabromide• sulfur trioxide• diphosphorus trioxide

• BH3

• CH4

• SiH4

• NH3

• PH3

• borane

• methane*

• silane

• ammonia*

• phosphine

COMMON NAMES, BINARY MOLECULES ENDING IN H

Group Work 2.6: MOLECULAR NOMENCLATURE

• SiCl4• SbF5

• P2O5

• BF3

• AsBr3

• SeO2

• N2O3

• Watch spelling of

unfamiliar elements!