5.1: Properties and Changes pg. 175
Chemistry is the study of the substances around us: what is in them, what they do, and what they are used for.
An understanding of chemistry teaches us how to change substances into new and useful products.
Physical and Chemical Properties:
Physical Property : a description of a substance that does not involve forming a new substance; for example, colour, texture, density, smell, solubility, taste, melting point, and physical state.
Chemical Property : a description of what a substance does as it changes into one or more new substances.
Physical and Chemical Changes:
Physical Change : a change that does not produce a new substance. Changes of state, a solid, liquid, and gas are examples of physical changes. Many physical changes can be reversed.
Chemical Changes : a change that produces a new substance(s).
If the products of a chemical reaction are not the same as the starting reactants, then a chemical change has taken place.
Indicators that a chemical change has taken place are:
1.
A new colour is produced
2.
Heat or Light is produced or absorbed
3.
Bubbles of gas are formed
4.
A solid material ( a precipitate) forms in a liquid
5.
The change is difficult to reverse

5.2: Identifying Physical and Chemical Changes pg. 180
A chemical property describes the ability of a substance to react to form something new. When this happens, we call it a chemical change.
In this activity, you will observe a number of physical and chemical changes. As you classify these changes, you will identify specific evidence of a chemical change.
Purpose: to collect and use evidence to identify physical and chemical changes.
Hypothesis:
Materials: eye protection
Bunsen burner lab apron spark lighter
2 test tubes test-tube stopper utility stand with clamp laboratory scoop warm water bath thermometer distilled water
CuSO
4
HCl NaOH
Magnesium ribbon wooden splint steel wool Lauric Acid
Cautions:
Hydrochloric acid and sodium hydroxide are corrosive. Sodium hydroxide can cause blindness if splashed in the eyes
Cooper (II) sulfate is toxic and an irritant. Avoid skin and eye contact. Wash any spills on the skin, in the eyes, or on clothing immediately with cold water. Report any spills to your teacher.
Use caution around the hotplate and water bath. Do not touch surfaces that might be hot.
This activity involves open flames. Long hair should be tied back and loose clothing tucked in.
Procedures:
1.
Prepare a data table in which to record your observations during the activity.
2.
Put on your eye protection and lab apron.

Part A:
Change 1:
1.
Add hydrochloric acid to a test tube to a depth of about 2 cm.
2.
Add two 1 cm strips of magnesium ribbon to the test tube. Check for evidence of change occurring. Test the bottom of the test tube with your hand for temperature changes. Record your observations.
3.
Place the test tube in a test tube rack and wait 30 s. for the gas produced to push any air out of the test tube.
Change 2:
1.
Clamp the Bunsen burner to the utility stand for stability.
2.
Light the Bunsen burner with a spark lighter; then light a splint from the burner flame.
3.
Hold the burning splint near the mouth of the “acid + magnesium” test tube.
Record your observations.
4.
Dispose of the contents of the test tube as directed by your teacher.
5.
Rinse the test tube with tap water.
Part B:
Change 3:
1.
Pour distilled water into a test tube to a depth of about 3 cm.
2.
Add about 0.5 g of copper (II) sulfate (as much as about half an Aspirin tablet) to the test tube.
3.
Stopper and invert the test tube several times to mix its contents well. Record your observations.
Change 4:
1.
Remove the stopper. Add a small ball of steel wool (about the size of an
Aspirin tablet) to the test tube (fig. 1)
2.
Stopper the test tube again and mix.
3.
Allow the solids in the test tube to settle to the bottom. Record your observations.

Change 5:
1.
Remove the stopper. Add about 5 drops of sodium hydroxide solution to the test tube.
2.
Slowly add drops of hydrochloric acid to the test tube. Gently swirl the test tube after every couple of drops. Continue adding drop wise until the solid disappears. Record your observations.
3.
Dispose of the contents of the test tube as directed by your teacher.
Part C:
Change 6:
1.
Examine a test tube of Lauric Acid
2.
Place the test tube in a warm water bath. Wait until the substance in the test tube completely liquefies.
3.
Remove the test tube and cool it in a stream of tap water until the contents solidify again. Record your observations.
Analyze and Evaluate: a) Classify each of the changes that you observed as either chemical or physical.
Use specific evidence from your observation table to justify your inference. b) Which changes were the most difficult to classify? Why? c) Give one example, from your everyday life, of a physical change that is; i.
Reversible. Justify your inference. ii.
Not reversible. Justify your inference.
Apply and Extend: d) Identify one chemical change in this activity that was reversible. What chemical could you add to reverse the change again? e) In change 2, you may have heard a “pop” when the burning splint was inserted into the mouth of the test tube. Name the gas produced in the test tube.

5.4: Patterns and the Periodic Table pg. 184
Elements : are pure substances that cannot be broken down into simpler substances.
The Periodic Table is used by scientists to explain and predict the properties of elements.
Chemical Periods and Groups
Period : each row of elements on the periodic table.
Group : each column of elements with similar properties.

Four Important Groups of Elements on the Periodic Table.
Group 1: Alkali metals
properties, soft, highly reactive metals
Group 2: Alkaline Earth Metals
these are light and reactive metals.
Group 17: Halogens
very reactive elements
Group 18: Noble Gases
these elements are stable, complete octet, rarely react.
Table 1: Summary of Properties of Metals and Non-Metals

Atomic Structure
Each element behaves different from each other. This is because their structures are different. The atom consists of a nucleus and orbits (energy shells).
99 % of the atom’s mass is located in the nucleus, small volume
(less then 1%), very dense and positively charged.
99% of the atom’s volume is located in the orbits, but its mass is less the 1%, and is negatively charged.
Subatomic Particles
Table 2: pg. 185
Atomic Mass: is the total of protons and neutrons found in the nucleus of the atom.
Atomic Number: the number of protons in the nucleus is called the atomic number of the element.

Electron Arrangements and the Bohr-Rutherford Model
Bohr-Rutherford diagram is a useful way to represent the arrangement of electrons around a nucleus for the first 20 elements.
Each orbit is shown as a ring around the nucleus. The first ring can only hold 2 electrons. The second and third ring can contain up to maximum of 8 electrons each.
Since there an equal number of electrons and protons the atom will be neutral.
Electron Arrangements and Reactivity
Noble gases have a stable arrangement, therefore; they do not react with other elements. The reason they are stable is they have a full complement of electrons in their outer orbit.
Fig. 6, Helium (He – 4), neon (Ne – 20), argon (Ar – 40) are stable because they have filled outer electron orbits.

Fig. 6, Lithium (LI -7), sodium (Na – 23), and potassium (K – 39), are reactive because of their single outer electron.
Other atoms are reactive because they do not have a full complement of electrons (octet) in their outer orbit. These are valence electrons.
Fig. 5, Bohr-Rutherford diagram of Sodium ( Na – 23)
Compounds are substances made up of two or more elements in a fixed ratio. It is the valence electrons of the atoms that allow this interaction.
Check Your Learning: 1 – 9, page 187.

5.5: Atoms and Ions pg. 188
Ion: a charged particle that results when an atom gains or loses one or more electrons
An Ion is created when an atom loses or gains an electron.
Sodium , found in group 1, has one valence electron. When it gives up this electron, it becomes an ion with a positive (+) charge. All atoms found group 1 has one valence electron, therefore they will also take on a positive 1+ charge.
Cation: is a positively charged ion.
The sodium atom loses its outermost electron to form an ion. The sodium ion is stable because its outer orbit is full, like that of neon
Anion: is a negatively charged ion.
Fluorine
The fluorine atom gains one electron to become a fluoride ion, F
-.
Fluoride is stable because its outer orbit is full, like that of neon.

Other Examples:
Aluminum
Sulfur
Hydrogen
The hydrogen atom is unusual. It has the ability to form a positive or negative ion. The hydrogen atom can achieve stability by
1gaining an electron, negative positive
1+
charge.
charge, or by losing it only electron,

Naming Ions
Ions are classified as either cation , positively charged or anion , negatively charged.
Naming ions, the positive ion keeps its element name, but the negative ion is identified by adding the “ ide
” to the stem of the atom‟s name. e.g.: Ox ygen forms the name ox ide .
Phosph orus forms the name phosph ide .
Check Your Learning, 1 – 9, pg. 191

5.6: Ionic Compounds pg. 192
Ions – an atom with a positive or negative charge due to a loss or gain of electrons
Ionic Compounds – a compound made up of oppositely charged ions
Electrolyte
– is a compound that separates into ions when it dissolves in water, producing a solution that conducts electricity.
Molecular Compounds – a compound that is not ionic, formed by the combination of two or more atoms held together with covalent bonds
Valence Electrons – the electrons in the outer shell of an atom, which determine its power to combine with other elements
Chemical Formula – a formula that uses symbols and numeral to represent the elements in a pure substance
Ionic Compounds
When atoms lose or gain electrons, they form ions. Ions take on a charge. If an atom loses one or more electrons, it will become positively charged.
Sodium (Na) has one valence electron, if it should lose its, sodium will have one more proton than electron and will take on a positive charge.
Proton
Electron
Charge
Sodium (1
11
+
10
-
1
+
+
) Chlorine (1
-
)
17
+
18
-
1
-

Chlorine (Cl) has 7 valence electrons, if it should gain one, chlorine will have one more electron than proton and will take on a negative charge.
Sodium + Chlorine → Sodium chloride
Calcium + 2 Fluorine → Calcium fluoride
Proton
Electron
Charge
Calcium (2
+
)
20
+
18
-
2
+
Fluorine (1
-
9
+
10
-
1
-
)
Calcium has two valence electrons, which it will give up, creating a positive charged ion (2
+
).
Fluorine has seven valence electrons, it requires to gain one electron, creating a negatively charged ion (1
-
).
One calcium ion requires two fluorine atoms to form a compound of calcium fluoride.

NaCl H
2
O
Figure 3: a) Under the microscope Sodium chloride appears as cubes. b) A crystal of sodium chloride could contain billions of alternating sodium and chloride ions. However, the number of sodium ions is always equal to the number of chloride ions, so their ration is 1:1.
Some ionic compounds are soluble in water. They dissolve and separate into individual ions. Water molecules surround each of these ions, preventing them from rejoining and forming the crystals again.
Figure 4: When ionic substances dissolve, their positive and negative ions are pulled away from the crystal by water molecules. The water molecules arrange themselves around ions in particular ways; the oxygen of water molecules are attracted to positive ions and hydrogen atoms are attracted to negative ions.

Aluminum can also react with chlorine gas. Each aluminum atom has three valence electrons to lose. Chlorine can only gain one electron. Each aluminum atom reacts with three chlorine atoms.
The resulting ionic compound is called Aluminum Chloride. When this molecule is dissolved in water, the aluminum atoms separate from the chloride atoms.

Properties of Ionic Compounds
Electrolyte
– is a compound that separates into ions when it dissolves in water, producing a solution that conducts electricity.
Most ionic compounds are electrolytes. When they are dissolved in water they form ions again. The presence of ions in water improves the electrical conductivity
Figure 6: Sodium chloride is an electrolyte because it separates into ions when it dissolves. A solution can conduct electricity only if it contains ions that are free to move.
Check Your Learning: Questions 1 – 13, pg. 195

5.7: Names and Formulas of Ionic Compounds pg. 196
Today, the number of known chemicals has grown to over 10 million. To keep track of them all, chemists have developed a systematic method of naming chemicals. The International Union of Pure and Applied Chemistry (IUPAC) is the organization that decides how chemicals will be named.
Naming Ionic Compounds
Many ionic compounds are made up of two different elements; consisting of a metal and a non-metal. Therefore the name of the compound will consist of two names. The first part will be the metal and the second the non-metal. The metal name will be named after the neutral atom of the metal and the second name will be the non-metal, but the ending will be changed to end with “ ide
”.
Chemical Formulas and Composition of Compounds
A chemical formula is a combination of symbols that represent a particular compound. e.g., Magnesium chloride MgCl
2
Magnesium ion Mg
2+
is combined with Chloride ion Cl
-
Ratio of atoms: one magnesium and two chloride atoms. This ratio is based on their valence electrons that are lost or gained.

Ionic compounds dissolve in water to form solutions that are able to conduct electricity, because they are made up of charged particles. The ionic compounds separate into their individual ions and are able to carry electric current.
Compounds that are not ionic, do not disassociate in water, therefore are not able to carry an electric current. Glucose is a
Molecular Compound , made up of non-metals, and are not electrolytes.
Table 1: Examples of Naming Ionic Compounds
Metal Metal ion Non-metal Non-metal ion Compound
Magnesium Magnesium ion Chlorine Chloride ion Magnesium chloride
Aluminum Aluminum ion Oxygen Oxide ion Aluminum oxide
Table 2: Names and Charges of Common Anions
Name of Element Name of ion fluorine chlorine bromine iodine oxygen sulfur nitrogen phosphorous fluoruide ion chloride ion bromide ion iodide ion oxide ion sulfide ion nitride ion phosphide ion
Ionic Charge
-
1
-
1
-
1
-
1
-
2
-
2
-
3
-
3
Ionic Symbol
F
-
Cl
-
Br
-
I
-
O
2-
S
2-
N
3-
P
3-

Writing Chemical Formulas of Ionic Compounds pg. 197
How can we name compounds?
How can we write formulas to represent these compounds?
There are over 100 elements identified on the Periodic Table .
These elements can be combined to make up thousands of different compounds.
From the periodic table, you are able to determine the ionic charges of elements, and then form ionic compounds. The ability to determine an element‟s ionic charge is dependent on the number of Valence electrons that element has.
Metals and Non-metals will combine to form ionic compounds, when the valence electrons are lost or gained.
Atoms that lose electrons take on a positive charge equal to the number of valence electrons lost by that element.
Sodium has one valence electron. It will lose that electron and take on a positive charge . There is one more proton in the nucleus than electrons found in the outer energy level.
Atoms that gain electrons take on a negative charge equal to the number of electrons gained to create an octet energy level.
Chloride has 7 valence electrons. It will gain one electron to create a complete stable octet. There is one less protons in the nucleus than electrons found in the outer energy level creating a negative charge.
Na
+
+ Cl
-
→ NaCl
Sodium ion + Chlorine ion → Sodium chloride

Writing Formulas for Ionic Compounds
A logical approach for writing formulas for ionic compounds involves a series of steps;
Step 1: Write the symbols, with the metal (element with the positive ionic charge) first.
Step 2: Write the ionic charge above each symbol to indicate the stable ion that each element forms.
Step 3: Determine how many ions of each type you need so that the total ionic charge is neutral.
Step 4: Write the formula using subscripts to indicate the number of ions of each type.
A second method is the
“CRISSCROSS”
method.
Crisscross Rule: Write the ionic charge above the symbols. Then crisscross the numbers, using them as subscripts.
Examples:
1.
Calcium and Iodine
2.
Aluminum and Sulfur
3.
Nickel and Oxygen

Naming Ionic Compounds
Naming ionic compounds is straightforward. Name the metal first and then the non-metal . The metal name is the same as the element name. The non-metal name is changed. The ending is changed to „ ide
‟. i.e., Calcium and Iodine (element names)
Compound Name: Calcium Iodide
Aluminum and Sulfur (element names)
Compound Name: Aluminum Sulfide

Elements with Multiple Ionic Charges pg. 198
Names and Formulas for Atoms with More than One Ionic
Charge
There are some metals that are able to form more than one type of ion. These metals can form two different types of ions, such as;
+ copper. Copper can have an ionic charge of positive one (1 positive 2 (2
+
).
) or
When naming these compounds, same rules apply, but Roman numerals in brackets are used to indicate the ionic charge of the metal. i.e., Copper and Chlorine
CuCl Copper (I) chloride
CuCl
2
Copper (II) chloride

Table 3: Names and Multiple Ionic Charges for Common Metals
Metal Chemical symbol of element
Chemical symbols of ions
Names of ions copper Cu
Cu
+
Cu
2+ copper(I) copper(II) iron Fe
Fe
2+
Fe
3+ iron(II) iron(III) lead Pb
Pb
2+
Pb
4+ lead(II) lead(IV) manganese Mn
Mn
2+
Mn
4+ manganese(II) manganese(IV) tin Sn
Sn
2+
Sn
4+ tin(II) tin(IV)

5.9: Polyatomic Ions pg. 202
Sodium compounds in figure 2 (sodium phosphate, sodium nitrate, and sodium erythrobate) are ionic compounds similar to the others you have learned about in this chapter. Polyatomic ions have similar characteristics, white solid, relatively stable, and an electrolyte, ionic compounds. For sodium phosphate, the cation is sodium, but the anion is phosphate (PO
4
-3
), which is a polyatomic ion.
A polyatomic ion is an ion that consists of a stable group of several atoms acting together as a single charged particle. The ionic charge of a polyatomic is shared over the entire ion rather than being just one atom.
Figure 2: The phosphate ion is made up of four oxygen atoms bonded to a central phosphorous atom.

Table 1 Formulas and Charges of Common Polyatomic Ions
pg. 202
3
2
3
3
3
2
4
2
4
3
4
Writing Formulas for Polyatomic Compounds
Rule 1: Write the symbols of the metal and of the polyatomic group.
Rule 2: Write the ionic charges.
Rule 3: Choose the number of ions to balance the charge.
Rule 4: Write the formula using subscripts.

Crisscross Method
Crisscross rule: Write the ionic charges above the symbols and crisscross them.
Naming Polyatomic Compounds
The naming of polyatomic compounds is relatively straightforward.
The name is simply a combination of the name of the Metal and the name of the Polyatomic ion. (Table 2) i.e., Metal , potassium (K
+
), and Polyatomic ion , carbonate (CO
3
2-
)
K
+
+ CO
3
2-
→ K
2
CO
3
Potassium + Carbonate Ion → Potassium Carbonate
The positive part of the compound is always written first in both the formula and the name.

Polyatomic ions make it possible to have an even wider range of ionic compounds, especially when they combine with metals that may have more than one ionic charge.
What is the formula for Lead (IV) carbonate ?

Lead has two different valences, but the Roman numeral tells you which one to use.
Rule 1: Write the symbols of the metal and of the polyatomic group.
Pb CO
3
Crisscross Rule: Write the ionic charges above the symbols and crisscross them.
Pb
4+
+ CO
3
2→ Pb
2
(CO
3
)
4
The formula Pb
2
(CO
3
)
4
, which must be reduces to Pb(CO
3
)
2
* Note the parentheses are included in the formula to show the number of CO
3
ions in the formula.

5.10: Molecules and Covalent Bonding pg. 206
NaCl H
2
O
Ionic Compounds form ions when dissolved in water. Metals with one, two, or three electrons in their outer shell lose electrons to non-metals, forming positive ions. Non-metals, have five, six, and seven electrons in their outer shell gain electrons, forming negative ions. Therefore the solution is able to conduct electricity.
If a solution does not conduct electricity, it must be made up of a different kind of compound.
Most compounds encountered every day, do not contain ions.
These compounds contain neutral groups of atoms, such as hydrogen and oxygen, share electrons to form stable arrangements.
Water ( H
2
O ) and carbon dioxide ( CO
2
), are two such molecular compounds. (solid, liquid, or gas)

Figure 6: Molecular models of a a) water, b) ammonia, c) nitric oxide
Hydrogen gas is a molecule formed when two hydrogen atoms are combined. Each of the Hydrogen atoms has only one electron. To become stable, both would have to gain an electron, 2 electrons in its outer orbit. The only way both atoms could gain an electron is by sharing electrons.
When two atoms share electrons, it results in a Covalent Bond .
There are many non-metallic elements that exist as covalently bonded molecules. A Diatomic molecule, is a molecule that is made up of 2 atoms.

Table 1: Common Diatomic Elements
Name of element Chemical
Symbol hydrogen H page 208
Formula of molecule
State at room
Temperature oxygen fluorine bromine iodine nitrogen
O
F
Br
I
N
H
2
O
2
F
2
Br
2
I
2
N
2 gas gas gas liquid solid gas chlorine Cl Cl
2 gas
Writing Formulas for Molecular Compounds
The formulas of many molecular compounds can be predicted using a method similar to Ionic compounds. The number of electrons that non-metals share to become stable is a clue to how many covalent bonds an atom can form.
Combining Capacity – is the number of electrons an atom can gain, lose or share to form a compound.
Table 2: Combing Capacities of Non-metal Atoms
4
C
3
N
2
O
1
H
F
Si P
As
S
Se
Cl
I

Carbon has 4 valence electrons in its outer orbit. If it loss these 4 electrons, it would have the arrangement similar to Helium and be a positive ion. If Carbon gained 4 electrons, it would have a similar arrangement to Neon, and it would be a negative ion.
Carbon cannot form either of the two ions. Instead Carbon gains 4 electrons by sharing. Carbon has a bond capacity of 4, therefore it can form 4 bonds. When carbon shares its four electrons with
Hydrogen it forms Methane gas (CH
4
). Each Hydrogen atom shares its electron with the Carbon and Carbon shares one with the
Hydrogen creating a covalent bond. This sharing allows each atom to complete its most outer orbit, creating a stable electron arrangement.

You can use the combing capacity to write the formulas of molecular compounds.
Rule 1 : Write the symbols, with the left-hand element from table 2 first, with the combining capacities.
C
4
S
2
Rule 2 : Crisscross the combining capacities to produce subscripts.
C
4
S
2
The formula is: C
2
S
4
Rule 3 : Reduce the subscripts it possible.
C
2
S
4
→
C
1
S
2
Rule 4
: Any „1‟ subscript is not needed.
The formula is: CS
2
Naming Molecular Compounds
Table 2: Common Names of some Molecular Compounds
Common name
Chemical formula
Use/occurances
Water
Ammonia
Nitric oxide
H2O the most commonly available molecular compound on
Earth; the “universal solvent”
NH3 used in window cleaners and in the production of fertilizers
NO an air pollutant produced in the automobile engine when gasoline is burned

Naming molecular compounds, prefixes are usually used. The prefixes are used to count the number of atoms when the same two atoms form different compounds.
Carbon Di oxide Gas (CO
2
Carbon Mon oxide Gas (CO)
)
If the first element is singular, the prefix Mon(o) is not required.
Prefix Number of atoms Sample molecular compound mon(o)- 1 carbon monoxide, CO di- tri-
2
3 carbon dioxide, CO
2 sulfur trioxide, SO
3 tetra- penta-
4
5 carbon tetrachloride, CCl
4 phosphorus pentafluoride, PF
5
Table 3: Prefixes Used for Molecular Compounds
Some compounds are named similar to the ionic compounds, for example; hydrogen sulfide (H
2
S). While other have more common names, such as; water (H
2
O), ammonia (NH
3
) and hydrogen peroxide (H
2
O
2
).