Unit 1 Foundations of Chemistry
Part A. Atomic Structure and Nomenclature
1. Introduction to Chemistry
Chemistry is the study of the properties of materials and the changes that materials undergo. It is the
central science, leading to a fundamental understanding of other sciences and technologies. It is an
extremely practical science that greatly impacts our daily living:
 Improvements to health care
 Conservation of natural resources
 Protection of the environment
 Provision of our everyday needs for food, clothing and shelter
By using chemistry, we have discovered helpful pharmaceutical chemicals, increased food production
and developed plastics.
2. Classification of Matter
Matter is the physical material of the universe. We can classify matter in the following ways:
The tremendous variety of matter in our world is due to the combinations of only about 100 very basic
substances called elements. Chemistry attempts to understand the properties of matter in terms of
atoms.
Summary:
 All matter is composed of elements.
 All elements are composed of very small particles called atoms.
3. Atomic Structure
Atoms contain two main regions called the nucleus and the electron cloud. The nucleus is
composed of two subatomic particles called protons and neutrons. The electron cloud is
composed of one subatomic particle called the electron.
The Bohr Model
The Electron Cloud Model
4. Pure Substances
A pure substance is matter that has distinct properties and a composition that doesn't vary from
sample to sample. Examples: Water, Oxygen and Table salt.
All substances are either elements or compounds. Elements are composed of only one kind of atom
(ex: O, H, Fe). Compounds are substances composed of two or more elements such as water and
table salt. A molecule is essentially a group of atoms bonded together. Note: Molecules of elements
consist of two or more similar atoms. Molecules of compounds consist of two or more different atoms.
Brown et al; Chemistry: The Central Science; p. 8
5. The Periodic Table
Periodic Table - A structured arrangement of elements that allows us to explain and predict physical
and chemical properties.
Symbols on the Periodic Table
6. Charges
The charges of the three key subatomic particles:
 Protons - Positively charged. Denoted: p+
 Neutrons - Neutrally charged. Denoted: n0
 Electrons - Negatively charged. Denoted: eElements on the Periodic Table are in electronically neutral form; i.e. they carry no charge. To be
electronically neutral, an atom must have an equal number of protons and electrons.
Example: Which atom below is electronically neutral?
7. Ions
The nucleus of an atom is unchanged by chemical processes, but atoms can readily gain or lose
electrons.
An ion is formed when electrons are removed from or added to a neutral atom. An anion is a
negatively charged ion. A cation is a positively charged ion. Example: How many electrons does Na +
have?
General rule: Metal atoms tend to lose electrons to form cations, whereas non-metal atoms tend to
gain electrons to form anions.
The red line divides metals from non-metals. It also separates cations from anions.
(Brown et al; p. 56)
How Elements Form Compounds
Much of chemical activity involves the transfer of electrons from one substance to another. Ionic
compounds are generally combinations of metals and non-metals. These compounds are held
together by attractions between opposite charged ions (like a magnetic attraction). This is called an
ionic bond. Molecular compounds are generally composed of non-metals only. Electrons are
shared between atoms. This is called a covalent bond.
Examples:
 Ionic bond – MgCl2 (magnesium chloride)
 Covalent bond – H2O (water)
Homonuclear (Diatomic) Molecules
These are molecules composed of only one type of element. The most common are:
P4 S8 H2 O2 F2 Br2 I2 N2 Cl2
Acronym: PS HOFBrINCl
H2 O2 F2 Br2 I2 N2 Cl2 are referred to as diatomic molecules since they are formed by pairs of
elements
Why would two chlorine atoms share a pair of electrons to form a covalent bond?
Polyatomic Ions
Polyatomic ions - groups of atoms that tend to stay together and carry an overall ionic charge.
Examples:
 nitrate – NO 3
 hydroxide – OH–
 bicarbonate – HCO 3

chlorate – ClO 3


sulfate – SO 24
phosphate – PO 34



8. Ionic Compounds
Are formed when metals and non-metals bond
The non-metals steal the electrons away from the metals (metals lose electrons and nonmetals gain electrons)
The result is a compound that is electrically neutral because the sum of the positive ions
equals the sum of the negative ions.
Naming Ionic Compounds
Basic Naming
 The first element (the metal ion) in the compound does not change its name.
 The second element (the non-metal ion) drops its ending and adds “ide”
 Example: LiCl = lithium chloride
 Practice: Name the following ionic compounds
1. CaCl2 ________________________
4. Sr3P2
________________________
2. MgBr2 ________________________
5. AlCl3
________________________
3. NaF
6. ZnI2
________________________
________________________
Assign Assignment #1
Naming With Polyatomic Ions
 As above, the first element (the metal ion) in the compound does not change its name.
 Use a table for the names of polyatomic ions. Their names do not change.
 There is only one positive polyatomic ion, NH4. If it is paired with a non-metal, drop the ending
and add “ide”.
 Examples:
NaNO3 = sodium nitrate NH4Cl = ammonium chloride
(NH4)2SO4 = ammonium sulfate
 Practice: Name the following ionic compounds
1. KNO3
________________________
5. Cs2SO3
2. Ca(OH)2
________________________
6. Mg3(PO4)2 ________________________
3. CaCl2
________________________
7. LiClO
________________________
4. (NH4)2S
________________________
8. Al(CN)3
________________________
Assign Assignment #2
________________________
Naming Compounds With Metals That Make More Than One Charge
Classical System
 Devised in 1787 by Guyton de Morveau, a French chemist
 Used extensively by industry
 The metals are referred to by their Latin/Greek names; suffixes are used to identify the charge.
 The suffix –ic is used for ions with the higher charge.
 The suffix –ous is used for ions with the lower charge.
Examples:
FeO = ferrous oxide
Fe2O3 = ferric oxide
Metal
Ion-lower charge
Ion-higher
Practice:
charge
1. Pb(NO3)4 ________________
2. FeCl3
________________
Iron
Ferrous
Fe+2 Ferric
Fe+3
+
+2
3. Cu2SO4 ________________________
Copper
Cuprous
Cu
Cupric
Cu
+4
4. SnSe
_________________
Tin
Stannous
Stannic
Sn
+2
Sn
Lead
Plumbous
Pb+2 Plumbic
Pb+4
+1
Mercury
Mercurous
Hg
Mercuric
Hg+2
Stock System
 The most common and modern way of naming ionic compounds.
 Devised in 1919 by Alfred Stock, a Prussion chemist.
 Use Roman Numerals in parentheses to indicate which charge of the metal used. This is
ONLY used for metals that make more than one charge.
 To figure out which charge is used, you need to look at the charge of the anion and
mathematically figure out what charge the positive ion has in order for the compound to be
neutral.
 Examples:
FeO  O has a 2- charge. Fe can be a 3+, 2+ or 6+. A positive charge of 2+ cancels out the
negative charge of 2-. Therefore, the compound is called iron ( ) oxide.
PbCl4 Cl has a 1- charge. There are 4 Cl’s in the formula so this adds up to a total charge of
4-. Pb can be a 2+ or a 4+ charge. The 4+ will cancel out the 4-. Therefore, the compound is
called lead ( ) chloride.
Ni3(PO4)2  PO4 has a 3- charge. There are 2 of them in the formula so this adds up to
a 6- charge. Ni can be a 2+ or a 3+ charge. There are 3 in the formula and they have to
add up to 6+ to cancel out the 6- charge. If we had three Ni2+, this would add up to 6+.
Therefore, the compound is called Nickel ( ) Phosphate.

Practice: Name the following ionic compounds where the metals make more than one charge.
1. TiP
__________________________ 5. Sn(NO3)4 ___________________________
2. AuBr3 __________________________ 6. CuSO3
___________________________
3. Co2O3 __________________________ 7. V3(BO3)5 ___________________________
4. FeI6
__________________________ 8. Cr3N2
Assign Assignment #3
___________________________
Practice Assignment: Name the following compounds. (For metals that make more than one
charge, use the Stock System.)
1. K2CO3
_____________________________________
11. AlN
2. Ca(OH) 2
_____________________________________
12. Ba(NO2) 2 _____________________________________
3. FeCl3
_____________________________________
13. LiClO3
______________________________________
4. Al2(SO4) 3
_____________________________________
14. Mg(CN)2
______________________________________
5. CoPO4
_____________________________________
15. RaCr2O7 __________________________
6. Mg(NO3) 2
_____________________________________
16. NiCO2
7. Zn3P2
_____________________________________
17. Pb(ClO)2 __________________________
8. NH4F
_____________________________________
18. NaCl
9. Li2SO3
_____________________________________
19. AgMnO4 __________________________
10. BeO
_____________________________________
20. Ag2MnO4 _______________________________________
_____________________________________
__________________________
__________________________
Assign Assignment #4
Writing Ionic Formulas
Basic Formulas
 Positive ion ALWAYS written first, negative ion second.
 The sum of the charges should add up to zero; ions add on until a neutral molecule is formed!
 Short cut: You can crisscross charge values by writing the number behind and below the
opposite element. (Make sure you reduce to the lowest possible formula.)
 If only one atom is needed, you do not write the subscript 1.
 Charges of the same value just cancel out. (1+1/1-1, 2+2/2-2 etc).
 Examples:
Sodium chloride  Na1+, Cl1- = NaCl
(charges cancel)
Calcium chloride  Ca2+, Cl1- = CaCl2
(criss-cross numbers)

Practice: Write the formulas for the following compounds.
1. sodium bromide
_________________________
2. strontium nitride
_________________________
3. potassium phosphide
_________________________
4. magnesium nitride
_________________________
5. zinc iodide
_________________________
Assign Assignment #5
Compounds With Transition Metals That Make More Than One Charge
 Some metals can make more than one kind of ion.
 Roman numerals in brackets are used in names to indicate the charge of the ion, NOT how
many are in the formula.
 You can still crisscross charge values by writing the number behind and below the opposite
element. Reduce to the lowest possible formula.
 Examples:
Iron (III) oxide  Fe3+, O2- = Fe2O3
Iron (II) oxide  Fe2+, O2- = Fe2O2 reduces to FeO

Practice: Write the formulas for the following compounds.
1. cobalt (II) chloride
_______________________
2. nickel (III) oxide
_______________________
3. lead (IV) selenide
_______________________
4. tin (II) nitride
_______________________
5. iron (III) phosphide _______________________
Assign Assignment #6
Compounds With Polyatomic ions
 Polyatomic ions are ions that, as a group of atoms, hold a charge.
 Their names generally end in –ite or –ate. (Some exceptions are hydroxide, cyanide etc.)
 Use the tables provided for charges of polyatomic ions.
 You can still crisscross charge values by writing the number behind and below the element or
polyatomic ion, however:
 Make sure you put the whole polyatomic ion in brackets if you need more than one!
 DO NOT change the formula of the polyatomic ion. Ex) NO3 has a 1- charge. It stays as NO3.
Use the 1- charge for your formula. DO NOT remove the 3!!!!
 Examples:
Sodium carbonite  Na1+, CO22- = Na2CO2
Aluminum sulfate  Al3+, SO42- = Al2(SO4)3
(use brackets)

Practice: Write the formulas for the following compounds
1. sodium chlorate
_______________________
2. magnesium phosphate
_______________________
3. silver nitrate
_______________________
4. barium sulfite
_______________________
5. iron (II) nitrite
_______________________
Assign Assignment #7
Practice Assignment: Write formulas for the following compounds
1. sodium bromide
_______________________
2. ammonium fluoride
_______________________
3. calcium carbonate
_______________________
4. nickel (II) phosphate
_______________________
5. lithium sulfite
_______________________
6. zinc phosphide
_______________________
7. silver phosphite
_______________________
8. tin (II) sulfide
_______________________
9. potassium permanganate _______________________
10. lead (II) nitride
_______________________
11. cobalt (II) carbonate
_______________________
12. sodium hydroxide
_______________________
13. iron (III) chloride
_______________________
14. magnesium acetate
_______________________
15. lead (IV) chlorite
_______________________
Assign Assignment #8


9. Covalent Compounds
Are formed when 2 non-metals bond
Here, electrons are SHARED, not transferred as they are in ionic compounds.
Naming Covalent Compounds
 A prefix is used to indicate the number of each element in the compound. (The exception is we
don’t use mono- for the first element).
Prefix
Number
 Charges are not of concern in covalent compounds.
Mono1
 Drop the ending of the second element
Di2
and add –ide.
Tri3
 Examples:
Tetra4
CCl4  carbon tetrachloride
Penta5
N2O3  dinitrogen trioxide
Hexa6
 Practice: Name the following covalent compounds.
Hepta7
Octa8
1. PBr3 ____________________________________________________
Nona9
2. P2O3 ____________________________________________________
Deca10
3. CF4 ____________________________________________________
4. SO2
____________________________________________________
5. N2O
____________________________________________________
Assign Assignment #9





Writing Covalent Formulas
Completely ignore the charges of the non-metals!!!
The prefixes will tell you how many atoms of each element you need.
DO NOT reduce compounds to lowest form! Leave them as the name states.
Examples:
Nitrogen trihydride  NH3
Dinitrogen tetraoxide  N2O4
Practice: Write the formulas for the following covalent compounds.
1. tetraphosphorus triselenide
________________
2. disilicon hexabromide
________________
3. diarsenic trioxide
________________
4. selenium monosulfide
________________
5. carbon tetrabromide
________________
Assign Assignment #10

10. Acids
Are covalent compounds (electrons are SHARED) that give off hydrogen ions when dissolved
in water.
They ALWAYS start with H+ bonded to some negative ion (either a non-metal or a polyatomic
ion).
There are three kinds of acids:
H+ and a polyatomic
o Hydrogen bonding with an –ate ions
o Hydrogen bonding with an –ite ions
o Hydrogen bonding with an –ide ions
H+ and a non-metal
(most of the time…exceptions are OH-, CN-)

Naming Acids
If the anion ends in –ate, it is called ________ic acid.


Ex) Hydrogen Nitrate, HNO3  _________________________ acid.

If the anion ends in –ite, it is called _______ous acid.
Ex) Hydrogen Nitrite, HNO2  __________________________ acid.

If the anion ends in –ide, it is called hydro____ic acid.
Ex) Hydrogen Nitride, H3N  ___________________________ acid.

Practice: name the following acids.
1. HI
___________________
4. H3P
__________________________
2. HIO3
___________________
5. H3PO4
__________________________
3. HIO2
___________________
6. H3PO3
__________________________
Assign Assignment #11



Writing Acids
Use the same crisscross method as with ionic compounds. H1+ will ALWAYS be at the start of
your crisscross.
Examples:
Hydrosulfuric acid (H1+ and S2-)  _________________
Carbonic acid (H1+ and CO32-)
 _________________
Nitrous acid (H1+ and NO21-)
 _________________
Practice: write the formula for the following acids
1. chromic acid
____________________ 6. sulfurous acid
2. carbonous acid
____________________ 7. Hydrocyanic acid ___________________
3. hydrochloric acid
____________________ 8. Oxalic acid
___________________
4. sulfuric acid
____________________ 9. Chlorous acid
___________________
5. hydrosulfuric acid
____________________ 10. Boric acid
___________________
Assign Assignment #12 and 13
___________________