1. 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

2. 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, , excellent reading on all topics.
Saunders CD-ROM: animated film clips
Reading, viewing assignment
Lecture Topics, Kotz,

3. Atomic Structure, Atomic Number, Mass
All matter (anything that has mass and occupies
volume) 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.

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

5. Comparative Mass, Charge: Nuclear Particles

6. 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

7. THE “NUCLEAR” ATOM
Nucleus,
C atom .
Electron cloud

8. X ATOMIC SYMBOLS Each atom of an element can be represented by a
symbol that describes how many protons, neutrons
and electrons are contained in this basic unit:
Mass number, A:
total #, p + n A
Elemental
symbol X Z
Atomic Number, Z:
#p (equals #e)

9. The Mass Number, A
Because atoms (and the subatomic particles) are so tiny,
a relative mass scale was setup to describe atomic
weights 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 essentially
makes the mass of each proton and neutron 1.00 amu.
Accordingly, A, the mass number of the atom, represents
both the total number of nuclear particles (p + n) and the
approximate mass of the atom (in amu’s)

10. Z, The Atomic Number
All known elements are listed in the familiar “Periodic
Table of the Elements” in order of increasing atomic
number, found usually in the upper right hand corner
above each element’s symbol.
The atomic number represents the number of protons
in the nucleus of every atom of that element.
Since every atom is electrically neutral, the number of
protons in the nucleus represents the total positive
charge of the nucleus, which is exactly balanced by
the total number of negatively charged electrons
outside the nucleus.

11. Examples, Atomic symbols
Atomic Mass #, p + n A X
Element
Atomic #, p (or e) Z p
= 146n
19 - 9p = 10n 19 F
238 U 9 92
9p, 9e
92p, 92e

12. Group Work 2.1: Atomic Symbols

13. 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 given
element is not fixed, and several different atoms of
a given element are generally found, varying in
atomic mass due to differing numbers of neutrons.
The different atoms of a given element which vary by
neutron count and by mass are described as “isotopes”
of that element.

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

15. Atomic Mass, revisited:
The atomic mass of any isotope of an element can be
approximated by a simple addition of the number of
p’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 (as
found in your Periodic Table) is a weighted average
of all the isotopes.

16. The given atomic mass for any element will be
closest to the most abundant element in most
cases. Consider below the calculation of the
atomic mass of magnesium, amu:
General Method of Calculation:
Average Mass Element, amu =
(mass, isotope A X % abundance A)
+ (mass, isotope B X % abundance B)
+ (mass, isotope C X % abundance C)......

17. Calculation of average atomic mass of Magnesium:
( x .7899) + ( x .1000) + ( x .1101)
= (18.95) + (2.499) + (2.861) = amu

18. Note:
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:
Mass, proton: amu
Mass, neutron, amu
Mass, 24 Mg , amu 12
1.007 amu X 12 =
1.009 amu X 12=
24.192

19. 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 nucleus
of an atom, a mass loss (or “mass defect”) always
occurs. This loss is thought to be the result of
a matter to energy conversion which holds together
the 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 or

20. The Periodic Table Lists all known elements in order of increasing
atomic number, left to right and top to bottom
Arranged so that elements of similar chemical
properties fall into the same column (called a group
or family)
Each horizontal line(called a period) represents
elements of a complete range of chemical properties.

21. 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 elements
becoming increasingly non-metallic
The end of the period features an inert, unreactive
element found only in the gas state

22. 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”

23. metalloids
metals
Non metals
Noble gases

24. Main Group or “Representative Elements”
Noble gases
inert gases
rare gases
8A (0)
Alkaline earth
metals 2A
Halogens 7A
Alkali metals 1A (except H)

25. Transition Metals 3 columns 8B: more alike “across” than “down”
Non- naturally occurring, newly found in lab

26. The inner Transition Metals:
Found beneath the main body of the PT, but
belonging to Period 6 and 7:
#58
#71
“Lanthanides”
#90
#103
“Actinides”

27. solids
liquids
gases
Note the seven common elements found in nature as “diatomic molecules”

28. Molecules and Compounds, Chapter 3
Lecture Topics, Kotz,
Atoms of the elements come together to form molecules and compounds.
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 O F2 Cl2 Br2 I2

29. 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 CH 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- )

30. Formulas for Molecules and Compounds
All compounds and molecules are represented by a
formula which summarizes number of atoms of each
element present:
H2SO CH3CH2OH NaCl
C12H22O F Fe2(SO4) (NH4)2CO H2O
MOLECULAR: formula represents all atoms in formula bonded
together in a single unit called a molecule...
IONIC: formula represents the simplest ratio of positive to
negative ions found in a sample of the compound...

31. TYPES OF FORMULAS
CONDENSED FORMULAS: summation of all atoms in formula, alcohol: C2H6O
STRUCTURAL FORMULAS: clue to connectivity in compound, alcohol: CH3CH2OH

33. 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

34. We’ll name in this lesson two specific types:
“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
“binary molecules” (two non-metals in formula )
CO PCl3 SO AsF3
Acids (H written first) will be done later...

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

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

37. POSITIVE IONS: Cations
CATIONS: positively charged ions; monoatomic
cations are formed from metals which have LOST
one or more electrons in compound formation:
- 1 e
Na (11p, 11e) > Na+ (11p, 10e) (all Group 1A)
-2e
Ca (20p, 20e) > Ca2+ (20p, 18e) (all Group 2A)

38. Common Metals, Fixed vs Variable Charge
Metals which form single cation
Metals which form several cations

39. 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 cation
Al3+ aluminum cation Ag+ silver cation

40. “One Charge Only” 3+ 3+ 2+ 2+ 1+ 1+

41. 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) will
justify all charges, single or multiple; now we simply
must recognize “which is which”

42. Naming Cations: Metals with Variable Charges
When a metal is known to form several different
cations of different charges, then the name of the
cation must include a Roman Numeral indicating
the charge of the ion:
Fe2+ Iron(II) cation Cu+ Copper(I) cation
Fe3+ Iron(III) cation Cu2+ Copper(II) cation
Sn2+ Tin(II) cation Bi3+ Bismuth(III) cation

43. 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 is
given by group number...

44. NEGATIVE IONS: Anions Monoatomic ANIONS:
Single nonmetallic atoms which have gained one or more
electrons in a chemical reaction and become negatively
charged ions :
+3e
N (7p, 7e) > N3- (7p, 10e) (all Group 5A)
+2e
O(8p, 8e) > O2- (8p, 10e) (all Group 6A)
+1e
F (9p, 9e) > F1- (9p, 10e) (all Group 7A)

45. Monoatomic Anions: Name, Charge3- 2- 1-
Fluoride
Chloride
Bromide
Iodide
Hydride Nitride Oxide
Phosphide Sulfide
Selenide
“ide”

46. FORMING IONIC COMPOUNDS
Make sure charges balance; cross multiply when
cation and anion charges are different:
Na Cl NaCl
Na S Na2S
Na P Na3P
Ba2+ Cl BaCl2
Ba2+ S BaS
Ba2+ P Ba3P2

47. GROUP WORK 2.3: Form Compound, Name
CATION, ANION
Mg H -
Fe S2-
Al P3-
Cd F -
Mn I -
FORMULA, NAME

48. 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; the
ion names end in “ate” or “ite” as follows...

49. Key Polyatomic Anion Formers: Know these!
Permanganate (7B)
like Perchlorate
Br, I same
as Cl
Chromate (6B)
like Sulfate

50. Polyatomic Anions of C, 4A
Most common: CO carbonate
HCO hydrogen carbonate,
“bicarbonate”
Others: CH3CO2- acetate (“C2H3O2-”)
CN cyanide

51. Polyatomic Anions of N, P 5A
Nitrogen: NO nitrate
NO nitrite
(Remember also: NH4+, ammonium; N3-, nitride)
Phosphorus: PO phosphate
HPO hydrogen phosphate
H2PO dihydrogen phosphate
(Remember also: P3-, phosphide)

52. GROUP WORK 2.4 CATION, ANION Cr3+ CO32- Ni2+ CN- Zn2+ NO2- Bi3+ H2PO4-
Pb N3-
FORMULA, NAME

53. Polyatomic Anions of O, S (6A) Cr (6B)
Oxygen: OH hydroxide
Remember also: O oxide
Sulfur: SO sulfate
SO sulfite
HSO hydrogen sulfate
HSO hydrogen sulfite
Remember also: S2-, sulfide
Chromium: CrO chromate

54. “BASES” and “SALTS”
The “hydroxides” and “oxides” of the metallic elements
are referred to as “bases”; all other ionic combinations
are referred to as “salts”
Bases:
Mg(OH)2
NaOH
CaO
Fe(OH)3
Salts:
MgCl MgHSO4 MgCO3
Na3PO NaNO Na2SO3
Ca(NO3)2 Ca3N CaSO4
Fe(CN)2 Fe(CH3CO2)3 Fe(H2PO4)2

55. 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
ClO chlorite
ClO chlorate
ClO perchlorate
Remember also: Cl-, Chloride
Manganese: MnO permanganate

56. 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 their
elemental name to end in “ide”
Polyatomic anions (memorized, Table 3.1, p. 110) end in “ite” or “ate”...

57. GROUP WORK 2.5 FORMULA NH4ClO NAME Cd(BrO2)2 Co(IO3)3 Ba(ClO4)2 KMnO4
Ag2CrO4
NAME

58. 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!

59. Binary Molecular Nomenclature Method:
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)

60. Typical Nomenclature NO2 SF6 ICl5 N2O5 CBr4 SO3 P2O3 nitrogen dioxide
sulfur hexafluoride
iodine pentachloride
dinitrogen pentoxide
carbon tetrabromide
sulfur trioxide
diphosphorus trioxide

61. COMMON NAMES, BINARY MOLECULES ENDING IN H
BH3
CH4
SiH4
NH3
PH3
borane
methane*
silane
ammonia*
phosphine

62. Group Work 2.6: MOLECULAR NOMENCLATURE
SiCl4
SbF5
P2O5
BF3
AsBr3
SeO2
N2O3
Watch spelling of
unfamiliar elements!