UNIT 1 – MATTER AND CHEMICAL BONDING
Unit Expectations:
Page Ref
After completing this unit, the student should be able to:
1. Briefly describe the contributions of Democritus, Dalton, Thomson,
Note, 11-14
Rutherford, Chadwick and Bohr to the development of atomic theory.
2. State the similarities and differences for the electron, proton and neutron
12-16
with respect to their location in the atom, charge and mass.
3. a) Define the following: Atomic number, mass number, isotope and
15-18
radioisotope.
b) Given the atomic number and the mass number of an atom, state the
number of electrons, protons and neutrons contained in the atom.
4. a) Describe how line spectra of elements tend to support the Bohr model of
Note
the atom that has electrons at different energy levels.
b) Draw Bohr-Rutherford diagrams and write electron configurations for the
Note
first twenty elements.
5. a) Define the following: Periodic law(modern version), period, family/group,
Note
valence electrons, octet, atomic radius, ion, anion,
cation, ionization energy, electron affinity.
6. a) Indicate the location of the Alkali Metals, Alkaline Earth Metals, Transition Note, 24-25
Elements, Halogens and Noble Gases in the Periodic Table.
b) State which elements occur in nature as gases, liquids or solids.
Note
7. State the general periodic trends in atomic number, metallic character and
Note
valence electrons for a period and for a family of elements.
8. State and explain the general periodic trends in atomic radius, ionization
energy, electron affinity and electronegativity for a period and for a family of 31-38,Note
elements.
9. Predict the stable ion that an atom will form based on its position in the
Note, 407
Periodic Table and write an ionization equation to represent this process.
10. Compare the properties of atoms and/or ions. (e.g. atomic radius)
Note, 31-34,53
11. Relate the reactivity of elements to their positions in the Periodic Table.
Lab, 25
12. Define the following: Chemical bond, ionic bond, covalent bond,
36, 54-57, 60-62
electronegativity, isoelectronic, double covalent
bond, triple covalent bond, polar covalent bond.
13. Use electronegativity values to determine the type of bond between
Note, 61
two atoms.
14. Describe the physical properties of compounds with ionic or covalent
Lab, 76-81
bonds.
15. Show the three steps involved in the formation of an ionic compound
Note, 54-55
and draw its Lewis structure.
16. Draw Lewis structures for covalent compounds.
Note, 56-57
17. Define the following: valence (Note), binary compound (pg. 666),
polyatomic ion (pg. 58), tertiary compound (Note).
18. State the valences of the required common cations and anions.
Write chemical formulas for binary compounds, binary acids, bases,
gaseous elements, oxyacids, tertiary compounds, acid salts, peroxides
and hydrates.
Using the Stock/IUPAC method, name binary compounds, binary acids,
bases, gaseous elements, oxyacids, tertiary compounds, acid salts,
peroxides and hydrates.
19. Using the –ous/-ic method, name binary compounds, bases, tertiary
compounds, acid salts and hydrates.
20. Using the greek prefix method, name binary compounds.
1
Note
Note, 67-69,
72
Note,64,69,
70,78
Note
Note, 71
THE DEVELOPMENT OF ATOMIC THEORY
Democritus –
Dalton (1803) –
Thomson (1897) –
Rutherford (1906) –
http://www.youtube.com/watch?v=ZPTGflQ2OS8
Chadwick (1932) –
Bohr (1913) –
2
Bohr-Rutherford Diagrams and Electron Configuration
Atomic Notation
Electron Configuration
Modern Definition of an Atom:
Atoms are composed of subatomic particles.
Subatomic Particle
Charge Symbol
Mass (in g)
Electron
Proton
Neutron
3
Radius (in m)
EXERCISE-ATOMIC STRUCTURE
Do Practice Problem 1 – 3 (pg. 14); Section Rev. Qu. 1-4 (pg. 21).
Also, do the following questions:
1. Indicate whether each of the following statements is true or false. If the statement is
false, rewrite it to make it correct.
a. Protons are negatively charged particles.
b. The mass of a proton is approximately equal to the mass of an electron.
c. The neutron is an uncharged particle found in the nucleus of atoms.
d. John Dalton discovered the electron.
e. In all atoms, the number of neutrons is equal to the number of protons.
f. Most of the mass of an atom is contained in the nucleus.
g. Isotopes of the same element have different atomic numbers.
2.
a. What is required for an electron in an atom to move from a lower energy level to a
higher energy level?
b) What is observed when the electron returns to a lower level?
3. Complete the following table:
Chemical Atomic
Mass
Number
Number
Number
Notation
Number
Number
of Protons
of Neutrons
Of Electrons
40
20 Ca
58
28
36
48
4. Draw Bohr-Rutherford diagrams for Elements 1 to 20 in the spaces in the chart on the
back of this page.
4
BOHR-RUTHERFORD DIAGRAMS
1
1H
4
2 He
7
3 Li
9
4 Be
11
5B
12
6C
14
7N
16
8O
19
9F
20
10 Ne
23
11Na
24
12 Mg
27
13 Al
28
14 Si
31
15 P
32
16 S
35
17 Cl
40
18 Ar
39
19 K
40
20 Ca
5
Name___________________
ISOTOPES AND AVERAGE ATOMIC MASS
RECALL
Protons:
Electrons:
e.g. 11𝐻
e.g.
16
8𝑂
Isotopes:
Isotopes Video
e.g. 168𝑂
18
8𝑂
17
8𝑂
Isotopes of an element have very similar chemical and physical properties because they
have the same numbers of protons and electrons. They have different masses because they
have different numbers of neutrons.
Radioisotope:
6
AVERAGE ATOMIC MASS & ISOTOPIC ABUNDANCE
The existence of isotopes can explain why the atomic mass on the periodic table is an average
atomic mass.
Average Atomic Mass
Mg consists of 78.70% Mg-24, 10.13% Mg-25, and 11.17% Mg-26. Calculate the
average atomic mass of Mg.
% Abundance
Lithium has two isotopes: Li-6 (6.00u) and Li-7 (7.00u). If the average atomic mass of Li
is 6.94 u, calculate the proportions (% abundance) of the isotopes.
7
ATOMS, ELEMENTS, AND THE PERIODIC TABLE
Dmitri Mendeleev (1834-1907)
The Modern Periodic Table
Periods:
Groups (aka Families):
Valence Electrons:
Periodic Table Movie
8
THE PERIODIC TABLE OF ELEMENTS
Alkali Metals
Noble Gases
Alkaline Earth Metals
Hydrogen
Halogens
Lanthanides and Actinides
9
PERIODIC TRENDS
When examining periodic trends, always look at:
As we go down a group:
As we go across a period (L to R):
IONIZATION EQUATIONS
-
e.g.
e.g.
1. IONIZATION ENERGY
e.g.
Which atom has the larger ionization energy: F or O? Why?
e.g.
Would it require more energy to remove an electron from a Ne atom or a F
ion?
e.g.
Which atom has the smallest ionization energy: Li or Rb? Explain.
10
IN GENERAL:
Down a Group:
Across a Period:
2. ATOMIC RADIUS
- e.g. Which atom would be larger: Be or Mg?
Explain.
v
e.g. Which would have the smallest radius: Mg or
Explain.
IN GENERAL:
Down a Group:
Across a Period:
3. ELECTRON AFFINITY
4. ELECTRONEGATIVITY
IN GENERAL:
Down a Group:
Across a Period:
11
Si?
THE VARIATION OF ATOMIC PROPERTIES
The purpose of this exercise is to demonstrate that there is a regular variation in the
properties of elements when the elements are arranged in order of increasing atomic
number. It also illustrates that predictions can be made based on this regular variation in
the properties of elements.
Procedure and Questions1. Using the data provided on the next page, plot a graph of atomic radius vs atomic
number. Put atomic radius on the vertical axis. Divide the vertical axis into intervals of
convenient size so that the largest atomic radius will appear near the top of the axis.
Put atomic number on the horizontal axis. Divide the horizontal axis into 36 intervals
using as much of the axis as possible. Join consecutive points with solid straight lines.
When the data for an element is missing, use a broken straight line to join the points for
the adjacent elements.
2. What would you estimate to be the missing value for the atomic radius of:
a. manganese?
b. selenium?
3. Would you expect the atomic radius for element 37 to be larger or smaller than the
atomic radius for element 36?
4. Using the same method as in part A, plot a graph of ionization energy vs atomic number.
5. What would you estimate to be the missing value for the ionization energy of:
a. manganese?
b. selenium?
6. What would be the approximate value for the ionization energy of element 37?
12
Atomic
Number
Element
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Atomic
Radius
(nm)
0.032
0.031
0.123
0.090
0.082
0.077
0.075
0.073
0.072
0.071
0.154
0.136
0.118
0.111
0.106
0.102
0.099
0.098
0.203
0.174
0.144
0.132
0.122
0.118
0.117
0.116
0.115
0.117
0.125
0.126
0.122
0.120
0.114
0.112
13
Ionization
Energy
(kJ)
1312
2372
520
899
801
1086
1402
1314
1681
2081
496
738
578
786
1012
1000
1251
1521
419
590
631
658
650
653
759
758
737
745
906
579
762
947
1140
1351
14
15
EXERCISE-TRENDS IN PROPERTIES IN THE PERIODIC TABLE
Also, do the following questions:
1.
a) Name the most metallic element in the periodic table.
b) Name the most nonmetallic element in the periodic table.
2.
Write the electron configurations of boron and aluminum. What is it about their
configurations that indicates that they belong in the same family?
3. Why does Be have a higher ionization energy than Li?
4.
a)Of the elements Al, Na and Cl, which one has the largest atomic radius? Why?
b) Of the elements B, Al and Ga, which one has the highest ionization energy? Why?
5. Which element has the largest electron affinity and why?
a) B, Li or F b) Br, Cl or I
6. Write an ionization equation for each of the following:
a) Na b) Al c) S d) Si
7) Which of the following are stable ions?
Ca+ , Cl+ , 0 , Cl , Mg2+ , K , H2
8) Which one of the following would have the smallest atomic radius?
02 , F , Ne , Na+ , Mg2+
9) Use the periodic table to predict which atom or ion in each of the following has the
greater radius:
a) sodium ion, aluminum ion
b) sodium ion, cesium ion
c) fluoride ion, iodide ion
d) sodium atom, chlorine atom
e) sodium ion, chloride ion
10) The atomic number of the hypothetical element, andromadine, is 117.
a) Which element group does andromadine belong to?
b) Compare the atomic radius and ionization energy of andromadine with the
atomic radius and ionization energy of radium and astatine.
16
CHEMICAL BONDS
Chemical bonds: the forces that attract atoms to each other in compounds.
PROPERTY
IONIC COMPOUNDS
COVALENT (MOLECULAR)
COMPOUNDS
Types of elements involved
What happens to electrons?
Type of bond that forms
State at Room Temperature
Solubility in Water
Electrical Conductivity in
Solution
BEFORE, classification of bond type was easy. All we did was look at the types of elements in
the compounds and that would tell us the bond type.
e.g. NaCl
e.g. CH4
Consider the following compound:
BeH2
Our previous experience and knowledge would seem to tell us that this is an ___________
compound.
However, the classification is incorrect! It is in fact a _____________ compound!
Therefore, our “inspection” method of classifying bond type is no longer sufficient to predict the
classification of bond type.
We must look to ELECTRONEGATIVITY for the answer!
Electronegativity: A measure of an element’s ability to attract electrons in a chemical bond.
By comparing the electronegativity values of two elements involved in a chemical bond, we can
determine whether they will transfer electrons (IONIC), or share electrons (COVALENT). For
elements that share electrons, we can also determine whether the sharing will be unequal
(POLAR COVALENT) or equal (COVALENT).
IONIC
POLAR COVALENT
17
TRUE COVALENT
e.g. Classify the type of bond present for each of the following compounds:
CH4
B2S3
CaO
H2O
IONIC COMPOUNDS – FORMATION & LEWIS STRUCTURES
Ionic compounds ____________ electrons to produce positive _________ and negative
____________. The oppositely charged ions ___________each other. The attraction is
what holds the ions together.
The attraction is __________. It is this attraction that gives ionic compounds HIGH
______________ and makes them __________ at room temperature. Since the ions are
___________, when they dissolve in water, they can _________________ by facilitating the
movement of the ___________ through the liquid.
RULES FOR DRAWING LEWIS STRUCTURES FOR IONIC COMPOUNDS
1. Write the ionization equation of each element to start.
Use this to determine the number of electrons transferred by each element.
Multiply the equations (if necessary) to equalize the number of electrons lost and gained.
2. Each ion goes in a set of square brackets [ ] with the charge outside the bracket
corresponding to the number of electrons lost or gained.
3. Alternate ions of opposite charges.
e.g. Sodium + Chlorine
e.g. Calcium + Bromine
e.g. Aluminum + oxygen
18
EXERCISE SHEET: IONIC AND COVALENT COMPOUNDS
Do Learning Check 13, 14, 15, 17 (pg. 79).
Also, do the following questions:
1) For each of the following, classify the bonding as being covalent, polar covalent or ionic:
a) BrCl
d) SiF4
g) CaO
j) NaI
m) MgCl2
b) CH4
e) Cs2S
h) OCl2
k) H2S
n) CCl4
c) NH3
f) Cl2
i) Na20
l) KBr
o) N2
2) Determine which one of the following compounds has the bonds with the greatest
polarity in it:
HF, CO2, FBr, H2O, F2O, HI
3) Arrange the following compounds in order of decreasing polarity of their bonds:
HBr, H2O, HF, CO2, HI, HCl, SI2
4) For each of the following, show the three steps involved in the formation of the ionic
compound and draw its Lewis structure:
a) lithium nitride
d) strontium chloride
g) magnesium sulphide
b) barium chloride
e) rubidium sulphide
h) potassium bromide
c) potassium oxide
f) calcium nitride
i) magnesium iodide
19
COVALENT (MOLECULAR) COMPOUNDS –
FORMATION & LEWIS STRUCTURES
RECALL:
Polar Covalent Bond:
Covalent Bond:
How can you tell if they will share equally?
e.g. Carbon dioxide
e.g. ammonia
e.g. carbon tetrafluoride
e.g. water
RULES FOR DRAWING LEWIS STRUCTURES FOR MOLECULAR COMPOUNDS:
1. Determine, from the chemical formula, the number of atoms of each type of element in
the compound. Use the last digit of the group number from the periodic table to
determine the number of valence electrons for each atom.
2. The element that requires the most electrons to complete its valence is called the central
atom. Draw the Lewis structure for this atom by placing one electron on each side of the
imaginary square enclosing the central atom before pairing up electrons.
a. If there is more than one of this type of atom, write them side by side and bond
them together using a line. Bonds form where there are unpaired electrons.
3. Use the unpaired electrons to bond additional atoms with unpaired electrons to the
central atom until a stable octet is obtained.
Draw the Lewis structures for the following compounds:
(a) H2O
(b) OF2
(c) CO2
(d) C2H2
NOTE: There are some atoms that can disobey the octet rule.
This occurs when they are bonded to highly electronegative atoms (e.g. F, Cl, Br).
When this happens, they can take more than 8e- (e.g. P, N, S) or less than 8e- (e.g. Be, B)
Be
4e- (2 bonds)
P& N
10 e-(5 bonds)
B
6e- (3 bonds)
S
12 e- (6 bonds)
e.g. BF3
SF6
NCl5
BeF2
20
EXERCISE SHEET: MOLECULAR STRUCTURE OF COVALENT
COMPOUNDS
Part A: Draw the Lewis Structure for each of the following:
(These molecules only have single bonds in them.)
1) Cl2
6) HCl
11) CHCl3
2) ICl
7) CH4
12) S2Cl2
3) H20
8) NF3
13) Br2
4) PH3
9) Cl20
14) SiBr4
5) H2S2
10) C2H6
15) PCl3
16) CBr4
17) P2H4
18) CF2Cl2
19) SeCl2
20) C3H8
Part B: Draw the Lewis Structure for each of the following:
(In this section, there is a mixture of all types of molecules i.e. molecules with single
bonds, molecules with double bonds, molecules with triple bonds and special case
molecules, too.)
1) CS2
9) N2H2
17) PCl5
24) C2H3Cl
2) HCN
10) BCl3
18) C2F4
25) SbCl5
3) BrCl
11) N2F4
19) I2
26) C4H10
4) C2F2
12) 02
20) C2Br2
27) C2FCl
5) HI
13) NCl3
21) OF2
28) CF4
6) CH20
14) H2S
22) SiCl4
29) C3H6
7) C2F6
15) SF6
23) XeF2
30) CH402
8) Si02
16) CH3Cl
21
VALENCE
Valence:
e.g. NaCl
e.g. HCl
e.g. N2
Some elements have more than one valence.
e.g. What is the valence of Pb in:
1. PbO
2. PbO2
Elements that have more than one valence are called multivalent elements.
Naming for Simple Binary Compounds
-
1.
e.g. NaCl
e.g. Na2O
e.g. CaH2
Writing Formulas for Simple Binary Compounds
1.
2.
3.
e.g. magnesium oxide
e.g. aluminum chloride
22
Naming Compounds with Multiple Valences
Some elements can have more than one valence.
e.g. Pb
This means it can form two compounds that are both called lead oxide. These compounds
have different chemical and physical properties.
Lead oxide =
To distinguish between them, we can use one of two methods:
1. IUPAC METHOD
2. Ous-Ic Method (Classical System)
e.g. PbO
e.g. PbO
e.g. PbO2
e.g. PbO2
Note: Often the historical name is used
Note: Roman Numerals are used
23
VALENCES
For each of the following compounds, state the valence of each element in the
compound:
eg.
Fe203
Fe is 3+, 0 is 2.
HBr
HF
H20
NH3
N203
CH4
SiCl4
C02
PCl3
PCl5
AsCl3
AsCl5
HI
AgI
Cu20
Cu0
Hg20
HgI2
SnF4
KI
Fe0
NaF
SnCl2
Sb203
Na2S
PbS
Pb02
MgS
ZnI2
CaBr2
Ba0
Al203
MnCl2
S02
S03
BCl3
SbCl5
SbCl3
LiF
BeI2
NiBr2
H2S
24
2
BINARY COMPOUNDS
PART A:
Name the following compounds.
NaCl ____________________________
Ca0 _____________________________
CaS _____________________________
H20 _____________________________
Na20 ____________________________
Mg3N2 __________________________
AlN _____________________________
Al203 ___________________________
PART B:
CaCl2 _____________________________
MgBr2 _____________________________
Ag2S ______________________________
AlI3 _______________________________
Al4C3 ______________________________
H2S _______________________________
SiC _______________________________
KBr _______________________________
Write the chemical formula for each of the following.
aluminum carbide _________________
aluminum oxide __________________
silicon carbide ___________________
sodium chloride __________________
magnesium bromide ______________
hydrogen sulphide ________________
aluminum iodide _________________
aluminum nitride _________________
hydrogen oxide ______________________
silver sulphide _______________________
calcium chloride _____________________
sodium oxide ________________________
calcium sulphide _____________________
magnesium nitride ___________________
calcium oxide _______________________
sodium hydride ______________________
25
3
BINARY COMPOUNDS (cont’d)
Write the chemical formula or the chemical name for each of the following using the
indicated method.
PART A: STOCK/IUPAC METHOD
Pb0 _______________________________ As2S5 ______________________________
Fe203 _____________________________ CuI2 ________________________________
Sn0 _______________________________ SbCl3 ______________________________
P203 ______________________________ Mn02 _______________________________
mercury(I) chloride ___________________ iron(II) oxide _________________________
antimony(III) iodide ___________________phosphorus(V) oxide __________________
tin(II) oxide _________________________ copper(II) bromide ____________________
PART B: -OUS/-IC METHOD
SnCl4 ___________________________
Sb203 ______________________________
CuBr2 ___________________________
FeBr2 ______________________________
As2S5 ___________________________
HgI ________________________________
mercurous chloride _________________
antimonous chloride ___________________
phosphoric sulphide
________________ ferric oxide __________________________
phosphorous oxide __________________ mercuric chloride _____________________
stannous bromide ___________________ cuprous iodide _______________________
phosphoric oxide ___________________ arsenic sulphide ______________________
stannic fluoride ____________________
ferrous oxide _________________________
PART C: GREEK PREFIX METHOD
S02 ________________________________As2S5 ______________________________
P2O3 ______________________________ N203 _______________________________
carbon dioxide _______________________carbon tetrachloride ___________________
phosphorus trichloride _________________carbon monoxide _____________________
diarsenic trioxide _____________________diphosphorus pentoxide ________________
antimony trichloride ___________________sulphur trioxide _______________________
sulphur hexafluoride ___________________diantimony trioxide ____________________
26
4
NAMING COVALENT COMPOUNDS
TWO NONMETALS:
DON`T WORRY ABOUT VALENCES!!
Use the prefixes to say how many atoms of each element are in the compound.
Mono
Di
Tri
Tetra
Penta
-
Hexa Hepta Octa Nona Deca –
e.g. carbon tetrachloride
e.g. H2O
e.g. NO
e.g. nitrogen trihydride
e.g. diphosphorus pentoxide
e.g. PCl3
Note: “mono” is never used on the first element.
27
One Type of
Element
One Atom
NOMENCLATURE
NOBLE
GASES
(NAMING)
Eg. He helium gas
Three or more
Elements
Two Types of
Elements
(Tertiary Compounds)
Two Atoms of
the same type
DIATOMIC
MOLECULES
I2 Br2 Cl2 F2 O2 N2 H2
I Bring Clay For Our New
House
Polyatomic Ions
can be Standard or
Derivative
(Binary Compounds)
Hydrogen +
Polyatomic Ion
Metal & nonmetal
(IONIC)
Metal is SINGLE
valence
IUPAC
Eg. MgO
Magnesium oxide
Metal is
MULTIVALENT
IUPAC Eg. Fe2O3
Iron (III) oxide
Ous-Ic Eg. FeO
ferrous oxide
Hydrogen &
nonmetal
Gas
IUPAC
Eg. HCl(g)
Hydrogen
chloride
Two Nonmetals
(COVALENT)
Aqueous
BINARY
ACID
Eg. HCl(aq)
Hydrochloric
acid
28
GREEK
PREFIX
Eg. P2O3
Diphosphorus
trioxide
NOTE
EXCEPTIONS!
N & P & Sb
OXYACIDS
Eg. HClO3(aq)
Chloric Acid
(Can use
standards or
derivative ions –
Per___ ate,
____ ite,
hypo____ ite)
Metal +
Polyatomic Ion
Metal is
SINGLE
valence
IUPAC
Eg. MgSO4
Magnesium
sulfate
Metal is
MULTIVALENT
IUPAC
Eg. FeSO3
Iron (II) sulfite
Ous-Ic
Eg. Fe3(PO4)2
ferrous phosphate
Naming Binary Acids, Bases and Gases
BINARY ACIDS
Recall:
- Acids
e.g.
HCl (aq)
HCl (g)
TO NAME BINARY ACIDS:
e.g.
HF(aq)
BASES
Recall:
-
hydro + ___________________ic acid
(name of non-metal)
e.g. H2S(aq)
Bases
e.g. NaOH
e.g. Fe(OH)2
e.g. ammonium hydroxide
e.g. copper (II) hydroxide
GASES
There are two types of gases.
Monoatomic (one atom) –
Diatomic (two atoms) –
29
BINARY COMPOUNDS, BINARY ACIDS, BASES AND
GASEOUS ELEMENTS
5
PART A – Name the following compounds using the IUPAC method.
CuCl ___________________________
Fe203 ________________________________________
SnBr2 ___________________________
As2S5 ___________________________
S02 _____________________________
Mn02 ___________________________
PART B – Name the following compounds using the –ous/-ic method.
HgCl2 ___________________________
FeCl3 ___________________________
Sb203 ___________________________
CuBr ___________________________
Fe0 _____________________________
As3N5 __________________________
PART C – Name the following compounds using the greek prefix method.
SiO2 ____________________________
C0 _____________________________
As2S3 ____________________________
SbI5 ____________________________
S03 ______________________________
CCl4 ____________________________
PART D – Write the chemical formula for each of the following compounds.
(Please note that the following section contains all of the methods of naming
that have been covered so far.)
mercuric bromide __________________
ferrous oxide _____________________
tin(IV) iodide _____________________
calcium hydroxide _________________
diantimony pentoxide _______________
hydroiodic acid ___________________
magnesium nitride _________________
silver sulphide ___________________
oxygen gas ______________________
hydrosulphuric acid _______________
zinc hydroxide ____________________
barium sulphide __________________
tin(II) oxide ______________________
carbon monoxide _________________
aluminum hydroxide _______________
30
BINARY COMPOUNDS, BINARY ACIDS, BASES AND
GASEOUS ELEMENTS (cont’d)
6
PART A – Name each of the following binary compounds using both the Stock and
-ous/-ic methods. In addition, name the last compound using the greek
prefix method, too.
Fe0 __________________________________________________________
Sb203 _________________________________________________________
Hg2S _________________________________________________________
SnCl4 _________________________________________________________
Fe203 _________________________________________________________
PBr3 __________________________________________________________
PART B - Write the chemical formula for each of the following compounds.
(Please note that the following section contains all of the methods of naming
that have been covered so far.)
sodium oxide ___________________
aluminum carbide ____________________
potassium iodide ________________
mercuric oxide ______________________
silicon oxide ____________________
aluminum hydroxide __________________
phosphorous iodide ______________
magnesium nitride ___________________
calcium bromide _________________
calcium hydroxide ___________________
sodium hydroxide ________________
hydrogen oxide _____________________
hydroiodic acid __________________
hydrofluoric acid ____________________
ferrous sulphide _________________
hydrogen fluoride ___________________
hydrosulphuric acid ______________
mercury(I) chloride __________________
iron(II) hydroxide ________________
magnesium hydroxide ________________
arsenic trifluoride ________________
nitrogen gas _______________________
antimony(V) sulphide _____________
helium gas ________________________
PART C - Name each of the following compounds. Remember that the Stock/IUPAC
method should always be used unless otherwise specified.
S03 ___________________________
As2S3 _____________________________
Hg0 ___________________________
HI(g) ______________________________
Ba(OH)2 _______________________
HBr(aq) ___________________________
Sn0 ___________________________
Ag2S ______________________________
SbCl5 __________________________
CuOH _____________________________
31
POLYATOMIC IONS AND RELATED OXYACIDS
Polyatomic Ion:
e.g. NO3-
There is a special way to name these ions. The following are the most common forms of the ions:
(See Valence Ion Reference Sheet)
THE STANDARDS:
ClO3-
Polyatomic Ions
chlorate ion
BrO3-
DERIVATIVE POLYATOMIC IONS
 ADD One O atom to a Standard:
“per_______ate ion”
e.g. ClO4-
IO3
NO3SO42-
Remove one O atom from a Standard:
“__________ite ion”
e.g. ClO2-
CO32-

PO43-
Remove two O atoms from a Standard:
“hypo______ite ion”
e.g. ClO-
 All standards have an “______ate” ending
THE STANDARDS:
HClO3
Oxyacids (Add H+)
chloric acid
HBrO3
DERIVATIVE OXYACIDS
 ADD One O atom to a Standard:
“per_______ic acid”
e.g. HClO4
HIO3
HNO3

H2SO4
H2CO3
Remove one O atom from a Standard:
“__________ous acid”
e.g. HClO2

H3PO4
Remove two O atoms from a Standard:
“hypo______ous acid”
 All standards have an “_____ic acid” ending
e.g. HClO
32
POLYATOMIC IONS AND RELATED OXYACIDS
PART A – Write the IUPAC name for each of the following polyatomic ions:
ClO3 ____________________________
PO43 ____________________________
BrO ____________________________
IO2 _____________________________
NO2 ____________________________
PO23 ____________________________
BrO4 ___________________________
ClO2 ____________________________
SO42 ___________________________
IO3 _____________________________
PART B – Write the chemical formula for each of the following polyatomic ions:
perchlorate _______________________
nitrate ____________________________
bromate _________________________
carbonate _________________________
periodate ________________________
phosphite _________________________
hypoiodite _______________________
sulphite ___________________________
bromite __________________________
hypochlorite _______________________
PART C – Write the IUPAC name for each of the following oxyacids:
HBrO3(aq) ________________________
HIO(aq) ___________________________
HClO4(aq) ________________________
H2CO3(aq) _________________________
HBrO2(aq) ________________________
HNO3(aq) __________________________
H3PO3(aq) ________________________
HClO(aq) __________________________
H2SO3(aq) ________________________
HIO4(aq) __________________________
PART D – Write the chemical formula for each of the following oxyacids:
chlorous acid _____________________
phosphoric acid _____________________
hypobromous acid _________________
sulphuric acid _______________________
chloric acid _______________________
nitrous acid ________________________
iodous acid _______________________
hypophosphorous acid _______________
perbromic acid ____________________
iodic acid __________________________
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7
NAMING TERTIARY COMPOUNDS
Tertiary Compound:
FOR A COMPLETE LIST, SEE VALENCE/ION REFERENCE SHEET
1.
2.
e.g. Li2CO3
e.g. FePO4
e.g. sodium hypobromite
e.g. lead (IV) perphosphate
e.g. NaClO
e.g. magnesium acetate
34
TERTIARY COMPOUNDS
8
Write the chemical formula for each of
the following:
silver phosphate ____________________
Write the Stock/IUPAC name for each of
the following:
Na2SO3 ______________________________
sodium sulphite _____________________
Ba(NO3)2 ____________________________
potassium perchlorate ________________
MnSO4 ______________________________
magnesium phosphite ________________
Hg2SO4 ______________________________
calcium sulphate ____________________
NaIO4 _______________________________
ferric sulphate ______________________
CaCO3 _______________________________
potassium nitrate ____________________
Fe2(SO4)3 ____________________________
magnesium hypobromite ______________
KNO3 ________________________________
zinc chlorite ________________________
NaClO3 ______________________________
calcium phosphate ___________________
Fe(NO3)2 _____________________________
sodium periodate ____________________
(NH4)2CO3 ____________________________
tin(IV) bromate ______________________
Mg3(PO4)2 ___________________________
manganese(IV) nitrate ________________
PbCO3 ______________________________
mercurous sulphite ___________________
Ca3(PO4)2 ____________________________
cupric perbromate ____________________
Fe(NO3)3 _____________________________
calcium carbonate ____________________
KI03 _________________________________
tin(IV) iodite _________________________
(NH4)3PO4 ____________________________
sodium phosphate ____________________
Al2(SO4)3 _____________________________
calcium hypochlorite __________________
MgSO3 ______________________________
sodium chlorite _______________________ Sn(NO2)2 _____________________________
magnesium perchlorate ________________
FeSO3 _______________________________
aluminum bromite _____________________ AlPO4 _______________________________
mercuric hypoiodite ___________________
NaBrO _______________________________
cupric sulphate _______________________ Na3PO3 ______________________________
35
FINAL PRACTICE: ALL NAMING
9
Write the chemical formula for each of
the following:
Write the Stock/IUPAC name for each of
the following:
antimonic perbromate _________________
AgClO3 ______________________________
aluminum iodite ______________________
HgBr0 _______________________________
ammonium carbonate _________________
K3PO4 ______________________________
iron(II) hypophosphite _________________
Cu2SO3 _____________________________
arsenic(III) hypochlorite ________________ HBrO2(aq) ____________________________
potassium nitrate _____________________
HgIO3 _______________________________
sodium sulphate ______________________ KCl02 _______________________________
calcium phosphite ____________________
Fe(BrO3)3 ____________________________
lead(II) periodate _____________________
MnO2 _______________________________
plumbic bromite ______________________
Hg(ClO3)2 ____________________________
cupric sulphite _______________________
KBrO4 _______________________________
sodium phosphate ____________________
Mg(IO2)2 _____________________________
potassium hypobromite ________________
MnCO3 ______________________________
magnesium perchlorate ________________
FePO2 _______________________________
barium iodate ________________________
AgClO _______________________________
antimonic bromate ____________________
NaNO3 ______________________________
arsenic(V) hypoiodite __________________
K2SO4 ______________________________
manganese(IV) chlorite ________________
Ba3(PO3)2 ___________________________
magnesium phosphate _________________ Sn(BrO)4 _____________________________
copper(I) sulphite _____________________ MgCO3 ______________________________
magnesium iodate ____________________
Ca3(PO4)2 ___________________________
lead(IV) hypochlorite __________________
Pb(IO)4 ______________________________
aluminum oxide ______________________
CuBrO4 ______________________________
manganous chlorate ___________________ CuSO4 ______________________________
silver perchlorate _____________________
CaCO3 ______________________________
36
UNIT 1 – REVIEW QUESTIONS
(Please note that the purpose of this sheet is to give you some practice questions to help you
prepare for the test. It does not have every type of question on it. The complete list of what
you are responsible for knowing for the Unit 1 test is on the Unit 1 Expectations sheet.)
Multiple Choice Practice:
pg 45 # 1,2,4,6 to 8
pg 89 #1, 3 to 8
pg 97 # 1 to 10
Short Answer Practice:
pg 45-46 #9,11,13,20,22,23
pg 89-90 #9 to 13, 16 to 24
pg 98 #11,12,15,16, 20 to 22, 31 to 33, 35b-d, 36, 40
Also Try these:
1. Complete the following table:
atomic
atomic
mass
symbol
number
number
40
20 Ca
number
of protons
number
of neutrons
58
number
of electrons
28
36
48
2. What is an isotope?
3. Draw Bohr-Rutherford diagrams for:
a) 147 N
b) 23
11Na
4. a)Which group of elements is most metallic?
b)Which group of nonmetals is most reactive?
c)Which group of nonmetals is least reactive?
5. How do the metallic properties of the elements in a period change as you move from left to right
across the periodic table?
6. a) What is the most reactive metal?
b) What is the most reactive nonmetal?
7. Why are the chemical properties of the elements in a given family quite similar, while those of the
elements in a given period are different from each other?
8. Which has the largest atomic radius Cl, Si, or Mg? Why?
9. Which has the largest ionization energy S, Se, or O? Why?
10. Which has the largest electron affinity S, Se, or 0? Why?
11. Write an ionization equation for:
a) Ca
b) N
c) Cl
d) B
12. State five properties of:
a) an ionically bonded substance
b) a covalently bonded substance
13. Is osmium a metal or a nonmetal?
14. Which of the following represent stable ions? F+ , Li , B3+ , O+ , S2 , Na2+ , N3
37
38
15. For each of the following, show the three steps involved in the formation of the ionic compound
and draw its Lewis structure:
a) potassium sulphide
b) beryllium fluoride
c) sodium nitride
d) cesium iodide
e) aluminum fluoride
f) magnesium nitride
16. For each of the following, classify the bonding as being nonpolar covalent, polar covalent or ionic:
a) I2
c) NCl3
e) SF6
g) Mg3N2
i) OF2
b) RbI
d) Ca3N2
f) SrBr2
h) AlF3
j) Cs2S
17. Consider the compounds beryllium oxide (Be0) and carbon monoxide (CO). One of the
compounds has a melting point of -215°C and the other has a melting point of 2530°C. Match
each compound with the melting point it is most likely to have.
18. Draw the Lewis Structure for each of the following covalent compounds:
a) HBr
f) C2Cl2F2
k) PBr5
b) F202
g) BCl3
l) SiH4
c) Br20
h) H2Se
m) CBr4
d) C2Cl2
i) SeH6
n) C2Br4
e) SiS2
j) C2Cl6
o)C3H4
39