Unit1: Matter and Chemical Bonding
Chemistry is the study of matter, its properties and its changes. There are many types of chemists: Analytical, Biochemical, Inorganic, Organic, Physical, Theoretical.
Matter
-has mass & occupies space
Pure Substance
-1 type of matter
-1 set of unique properties
Mixture
-2 or more pure substances
-separated by physical methods
Element Compound Mechanical Mixture Solution
-1 type of atom -2 or more atoms -2 or more phases -1 phase
-on periodic table chemically bonded -heterogeneous
- ex. -can be broken down - ex.
by chemical means
- ex.
-homogeneous
-uniform comp.
- ex.
Parts of the Atom
1. Proton- positive particle found in nucleus
2. Neutron- neutral particle found in nucleus
3. Electron- negative particle found orbiting the nucleus
Atomic number = number of protons
- Each element has a unique # of protons
- Appears on the periodic table
- In a neutral atom # protons = # electrons
Bohr-Rutherford Diagrams
Ex. Calcium
Lewis Structures (Lewis Dot Diagrams)
- Valence (outer orbit) electrons only
- e‘s surround atomic symbol ex.

Describing and Measuring Matter
Physical and Chemical Properties- help describe and identify matter.
Physical Properties- are observations about matter that can be measured (quantitative ex. melting point, mass) or that are descriptions (qualitative ex. colour, state)
Chemical Properties- are observed when one kind of matter is changed into a different kind of matter (ex. combustibility, reaction with acid)
Physical Change- a change in form without changing the chemical composition
Ex.
Chemical Change- a change in chemical composition producing a new substance
Ex.
Taking Measurements
- Always record all digits you are certain about and a final uncertain digit
- These recorded digits are called significant figures or significant digits
- The number of significant figures (sig figs) is important when performing calculations (see page 614)
Accuracy and Precision
Accuracy- how close your measurement is to the real value
Precision- exactness of a measurement or closeness of a series of measurements

1.1: The Nature of Atoms
Parts of the Atom / Average Atomic Mass
Atomic Symbols
A= mass number = # protons + # neutrons
Z= atomic number = # protons
# neutrons =
Isotopes – atoms of an element that have the same number of protons but different number of neutrons
Ex.
Average Atomic Mass
The mass of an atom is expressed in atomic mass units (µ). One atom of C-12 has a mass of exactly 12 by definition. The masses of all other atoms are measured in relation to C-12.
The isotopic abundance (relative amount each isotope is present in nature) needs to be taken into consideration when calculating average atomic mass.
The average atomic mass is the average of the masses of all the element’s isotopes.
Taking average of your course marks:

Taking average of assignments weighted differently:
Example: Finding the Average Atomic Mass of Magnesium
Mg-24 (23.985 µ) is 78.70 % of all Mg
Mg25 (24.985 µ) is 10.13 % of all Mg
Mg26 (25.983 µ) is 11.17 % of all Mg
Avg. Atomic Mass = sum of the mass of each isotope multiplied by its abundance
=
Try Ge
NB- Mass Spectrophotometer uses magnetic field to separate ions with the same charge but different masses to find the relative abundance of each isotope (read pg.
281)
Homework: pg. 19 #1-4, 10 pg. 21 #9

1.2: The Periodic Table
Periodicity
The periodic table is arranged according to the atomic number of the elements. This organization shows that the chemical and physical properties repeat in a regular, periodic pattern.
An element’s period number is the same as the number of energy levels that its electrons occupy.
For the main group elements, the number of valence e‘s is equal to the group number
(1,2) or equal to the group number minus 10 (13-18).
Categories of Elements in the Periodic Table
-groups of elements that have similar properties are given a family name

1.3: Explaining Periodic Trends
Atomic Radius (Size)
A n atom’s size is measured by its radius: the distance from the nucleus to the approximate outer boundary of the cloud-like region of its electrons
As you go down a group, the size of the atom increases. The valence e‘s are farther and farther away from a positive nucleus and the inner electrons shield the valence e-
‘s from the positive charge.
Across a period, the size of an atom decreases. Even though the number of electrons increases, they stay in the same energy level. The increasing positive charge of the nucleus pulls the electrons closer, reducing the atom’s size.
Ionization Energy
An atom that gains electrons becomes negatively charged (anion). An atom that loses electrons becomes positively charged (cation). The energy required to remove an e- from an atom is called ionization energy (IE).
Ionization energy _____________________ down a group and _______________ across a period.
Electron Affinity
Electron affinity (EA) is the change in energy that occurs when an e- is added to the outer energy level of an atom to form an anion. A high EA means that energy is released when an atom gains an e-, and that it is likely to form an anion. A low EA means that energy is absorbed when an atom gains an e-, and that it is not likely to form an anion.
Electron affinity generally ______________________ down a group and
____________________ across a period.
Homework: pg. 40 #1-5, 9,10

2.1: The Formation of Ionic and Covalent Bonds
Chemical Compounds
Elements consist of only one type of atom. Two or more atoms chemically bonded are called a compound. Compounds can be classified into two groups: ionic and covalent.
Chemical bonding involves the interactions between the valence electrons of atoms. In ionic bonding one atom loses its valence electron(s) and the other atom gains the electron(s). In covalent bonding the atoms share electrons. When an atom loses, gains or shares electrons through bonding to obtain a filled outer energy level, the resulting compound is stable.
A) Ionic Compounds
Most ionic compounds are crystalline solids with high melting points. They are usually highly soluble in water and conduct electricity.
Ionic bonding occurs between a metal and a non-metal. The metal transfers its electrons to the non-metal so that both atoms achieve a filled outer energy level. Most atoms have eight electrons in their filled outer energy level, but there are many exceptions.
Ex. 1 sodium chloride
Ex. 2
Solid Ionic compounds form a rigid lattice (see pg. 79) of interlocking sodium and chloride ions. The smallest unit of this structure (NaCl) is called a formula unit (not a molecule!).
Ionic compounds do not conduct electricity in the solid state as the particles are not free to carry the current. In the molten state the lattice structure is broken down and the ions are free to move and conduct electricity. When an ionic compound is dissolved in water the ions are free to move and the solution is a good conductor of electricity.

B) Covalent Compounds
A covalent bond consists of a pair of shared electrons. Two atoms of the same element can form a bond. These diatomic elements include:
Ex.
-count the electrons in a covalent bond as if they belong to each of the bonding elements.
Atoms can share two or three pairs of electrons.
Ex.
Covalent compounds have a wider variety of properties than ionic compounds. Most have low boiling points and do not conduct electricity in the solid, liquid, or gaseous state.
Covalent compounds do not break into ions when they melt or boil. Their atoms remain bonded together as molecules and hence covalent compounds can be called molecular compounds.
Ex.
When a covalent compound is in its liquid or solid state the molecules are held together by intermolecular forces.
C) Metallic Bonding
Metals do not form ionic bonds or covalent bonds with one another. The force that holds metal atoms together is called a metallic bond. Both pure metals and alloys contain metallic bonds. Alloys are homogeneous mixtures of two or more metals (ex. bronze = copper, tin, lead).

The properties of metals can be explained by their metallic bonds.
1) Metals are good conductors of electricity because e-s can travel to all cores as a new e- comes in from a current the valence e‘s can get popped off and out
2) Metals are good conductors of heat because the mobility of the electrons can transfer heat energy
3) Metals are shiny because light bounce off moving e‘s
4) Metals are malleable because there are no bonds to break
Electronegativity
The electronega tivity of an atom is a measure of that atom’s ability to attract electrons in a chemical bond.
Electronegativity ________________ down a group and ______________ across a period.
Ex1.
Ex.2.
ΔEN
The difference between electronegativities (Δ EN) helps us decide what type of bond is being formed. Always take the larger EN value and subtract the smaller EN value from it.
Covalent bonds can be divided into two categories:
1) Polar Covalent Bonds- form when the difference in EN is between 0.5 and 1.7.
This ΔEN is not great enough for the less electronegative atom to transfer its electrons to the other atom (ionic bonding) but great enough for the bonding electron pair to spend more time near the more electronegative atom.

Ex. water
The bond between O and H has a ΔEN of ____. The oxygen atom attracts the electrons more strongly than the H. The O obtains a slight negative charge and the H obtains as slight positive charge. The overall molecule has a dipole with the negative end near the O and the positive e nd near the H’s.
2) Non-Polar Covalent Bonds- form when the difference in electronegativity is between 0 and 0.5. This ΔEN results in equal sharing of the bonding electrons.
Ex. Methane
The bond between C and H has a ΔEN of ____. The atoms share the bonding electrons equally.
Molecular Shape
Electron arrangement affects the shape of molecules. Electron pairs involved in bonding are called bonding pairs. Electron pairs not involved in bonding are called lone pairs. Electrons arrange themselves around molecules so that they are as far apart as possible from each other.
A tetrahedron shape allows four e- pairs to be at a maximum distance from one another.
Lone pairs repel bonding pairs strongly and cause the shape of the tetrahedron to change.
Ex.1
Ex. 2
Homework: pg. 63 #2-6, 10a-d

2.2: Writing Names and Formulas for Ionic and Molecular Compounds
Naming Ionic Compounds
Recall:
- draw the Lewis structure for the reaction of calcium with chlorine
Chemical Formulas
The chemical formula tells us the elements that make up a compound and the number of each atom in the smallest unit of the compound. For ionic compounds the formula unit tells us the ratio of the two types of atoms.
Valence Numbers
In an ionic bond the valence is the charge on the ion that is formed. If an atom loses e-
‘s to become a cation, it is assigned a positive valence number. If the atom gains e-‘s to become negative, it is assigned a negative valence.
See periodic table in back of textbook
Ex.
Some larger elements can have more than one valence due to complex e distribution.
Ex.
Chemical formulas can be determined using Lewis structures or using valences. The sum of the valences must equal zero. The cross-over rule can be used:
1. write the symbol of the metal (cation) followed by the non-metal (anion) including their valences.
2. cross the charges down to become subscripts without the charges.
3. remove subscripts of 1 and reduce to lowest common number.

Polyatomic Ions
-groups of atoms that stay together to form an ion
Ex.
See Table 2.4 (must be memorized)
Chemical formulas for compounds containing polyatomic ions are determined the same way
Ex.
IUPAC Nomenclature
The International Union of Pure and Applied Chemistry established rules for naming compounds. There are still old-fashioned common names for some compounds (ex water, acetic acid).
For ionic compounds:
1. write the name of the less electronegative element (usually metal)
2. if the metal has more than one valence write the valence in brackets in Roman numerals
3. write the name of the more electronegative non-metal but change its ending to
“ide” ex.
Handout:
1. box of polyatomic ions on pg. 2
NB ate  ite : lose one O per added: add one O hypo & ite: lose 2 O rules:
See #2 on pg. 8

2. an older method of naming metals with more than one valence is to add the ending ic for larger valence and ous for smaller valence ex.
Nomenclature of Compounds that contain Hydrogen
3. A) IF hydrogen is the less electronegative element write the word “hydrogen” as the first part of the compound
Ex.
B) IF hydrogen is the anion change its name to “hydride”
Ex.
Acids are composed of hydrogen dissolved in water (given aq after chemical formula).
Binary acids (H + one element) are named by using the word
“hydro” + element + ic acid
Ex.
Acids that contain a polyatomic ion that ended in “ate” become
_________ ic acid
Ex.
_____________ ite becomes _____________ous acid
Ex. per ___________ate becomes per ___________ ic acid
Ex. hypo_____________ite becomes hypo _____________ ous acid
Ex.

1
2
3
4
5
6
Naming Covalent Compounds
A prefix is added to both non-metals to indicate the number of atoms each element in one molecule of the compound.
Number of atoms Prefix
7
8
9
10
The prefix is left out when there is only one atom of the first element in the name.
Ex.
Homework: pg. 70 #8-10,12; pg. 73 #1-10; pg.75 #5-10; worksheets

5)
6)
3)
4)
7)
1)
2)
18)
19)
20)
15)
16)
17)
8)
9)
10)
11)
12)
13)
14)
6)
7)
4)
5)
1)
2)
3)
Name the following compounds (IUPAC System Preferred)
Mg(OH)
BeF
2
KNO
3
2
_______________
_______________
_______________
21)
22)
23)
Ba
3
N
2
H
2
S
(aq)
_______________
_______________
BaCl
2
.
2 H
2
O _______________
NH
3
LiOH
Na
2
CO
3
HgO
HClO
Zn
BF
3
3
P
2
(aq)
AgHSO
3
Sn(ClO
4
)
2
_______________
_______________
_______________
_______________
_______________
_______________
_______________
_______________
_______________
HgCH
3
COO _______________
Mg(OH)
2
_______________
FeS
HF
(g)
CuCl
2
Na
2
S
HI
(aq)
Zn(OH)
2
_______________
_______________
_______________
_______________
_______________
_______________
24)
25)
26)
27)
28)
29)
30)
31)
32)
33)
34)
38)
39)
40)
35)
36)
37)
N
2
S
3
HNO
3
Ca(ClO
2
CO
Al(OH)
3
NaNO
3
)
2
BeO
Pb(HCO
3
)
2
FePO
4
_______________
_______________
_______________
_______________
_______________
_______________
_______________
_______________
_______________
CuSO
4.
5 H
2
O _______________
P
2
O
5
_______________
PbBr
2
Ag
2
O
SnO
2
Fe
2
O
3
Cu
3
P
2
Sn(ClO
4
)
4
_______________
_______________
_______________
_______________
_______________
_______________
CH
3
COOH
(aq)
_______________
HBr
Be(OH)
2
Sr(NO
3
)
2
As
Ni(NO
Na
2
2
(aq)
O
S
5
3
)
3
_______________
_______________
_______________
_______________
_______________
_______________
21)
22)
23)
24)
25)
26)
27)
CuCrO
4
Li
2
Li
2
Ca
3
O
SnO
2
P
2
NaH
CO
3
Al(OH)
3
_______________
_______________
_______________
_______________
_______________
_______________
_______________

11)
12)
13)
14)
7)
8)
9)
10)
2)
3)
1)
4)
5)
6)
8)
9)
10)
11)
12)
NaClO
Zn(MnO
(NH
NaNO
3
K
2
O
4
)
2
3
SO
4
)
4
2
_______________
_______________
_______________
_______________
_______________
28)
29)
30)
31)
32)
FeBr
K
3
3
PO
PbO
Sn
3
2
P
2
HClO
3
(aq)
_______________
_______________
_______________
_______________
_______________
13)
14)
15)
16)
17)
18)
BaSO
4
MgO
Sn(CN)
Fe
H
2
4
2
(CO
3
)
Pb
3
S
(aq)
(PO
5
)
4
3
_______________
_______________
_______________
_______________
_______________
_______________
33)
34)
35)
36)
37)
38)
19)
20)
Ca(OH)
2
AlF
3
_______________
_______________
39)
40)
Write the formula for each of the following compounds
HC1O
FeSO
3
3(aq)
MnH
2
HgNO
3
_______________
_______________
_______________
_______________
Sn(SO
5
)
2
_______________
CaSO
4
.
2 H
2
O _______________
As
H
2
(Cr
2
S
(g)
2
O
7
)
5
_______________
_______________ silver nitrate stannous chloride aluminum hydroxide
__________
__________
__________ ammonium dichromate __________ ammonium hydroxide __________ ammonium sulphate __________
25)
26)
27)
28)
29)
30) lead (ll) acetate carbon tetraiodide mercuric iodide copper (I) acetate plumbous iodite potassium acetate
__________
__________
__________
__________
__________
__________ perchloric acid iron (III) carbonate silver perchlorate mercury (II) chloride nitric acid zinc nitride boron sulphate potassium oxide
__________
__________
__________
__________
__________
__________
__________
__________
35)
36)
37)
38)
31)
32)
33)
34) manganese (III) chromate __________ calcium sulphite __________ tin (II) nitrate potassium sulfide
__________
__________ hydrochloric acid barium nitrate silver hydride aluminum oxide
__________
__________
__________
__________

20)
21)
22)
23)
24)
15)
16)
17)
18)
19) magnesium sulfide hydrogen iodide calcium fluoride potassium carbide mercuric oxide arsenous nitride copper (I) oxide nitrous acid ferric oxide lead (IV) nitrate
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
39)
40)
41)
42)
43)
44)
45)
46)
47)
48) calcium chloride potassium hydride sodium sulfide stannous fluoride plumbous oxide calcium phosphate ferric hydroxide copper (I) chloride hydrofluoric acid magnesium hydroxide
14)
15)
16)
17)
7)
8)
9)
5)
6)
10)
11)
12)
13)
Write the formula for each of the following compounds
1) hydrogen sulfide __________ 18)
2)
3)
4) ammonium nitrate stannic phosphate manganese (IV) nitrite
__________ 19)
__________ 20)
__________ 21) hydrogen acetate aluminum sulfide
__________
__________ hydrobromic acid __________ lead (II) hypocarbonite __________ ammonium sulfite silicon cyanide chromium (II) phosphide __________ 24) mercury (I) fluoride cupric sulfate calcium nitride zinc phosphide boron sulphate
__________ 22)
__________ 23)
__________ 25)
__________ 26)
__________ 27) potassium permanganate __________ 28)
__________ 29)
__________ 30) silver acetate phosphoric acid iron (III) cyanide ferrous bromate boron bromide magnesium nitrate silver hydride
__________
__________ magnesium dichromate __________ lithium carbonate __________
__________
__________
__________
__________
__________ rubidium oxide aluminum sulphate barium hydride manganese (II) oxide
__________ 31)
__________ 32)
__________ 33)
__________ 34) magnesium nitrate ferric permanganate silver sulphide nickel (II) arsenide
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________

Lewis Structures
Recall - use the last digit of group # from periodic table to determine the # valence electrons
- place 1 electron on each of 4 sides before pairing electrons
use unpaired electrons to bond additional atoms
Draw Lewis structures to illustrate the bonding between the following pairs of elements: a) Na and Cl b) Li and O c) d)
Mg and N
Ca and P
Draw Lewis structures to illustrate the bonding in the following molecules: a) HCl d) NH
3 b) CH
4 e) CCl
4 c) CO
2 f) SCl
2

ex. SF
4
2.3 Comparing the Properties of Ionic and Molecular Compounds
Extending the Lewis Theory of Bonding
-Sidgwick showed that one atom can contribute both electrons that are shared and that an octet of electrons is desirable but not necessary
Procedure for drawing Lewis Structures
1.
count total # valence electrons (add for anions, subtract for cations)
2.
place less electronegative atom in center and arrange other atoms symmetrically around the central atom
3.
place a bonding pair of electron between the central atom and each surrounding one
4.
fill the peripheral atoms with electrons (8 or 2)
5.
any remaining electron go to the central atom
6.
only give double or triple bonds if central atom does not have an octet and the bond does not cause the total # ex. NO
3
valence electrons to be > 8
- more than one Lewis structure possible (resonance structures)
- expanded valence energy levels are possible for some atoms due to their size (reason yet unknown)

Draw structures for: a) CF
4 b) BF
3 c) NF
3 d) OF
2 e) NF
5 f) SF
6
Valence Shell Electron Pair Repulsion (VSEPR) Theory allows us to predict the geometry of molecules based on the repulsion of electron pairs. There are five basic geometries based on there being a maximum of six bonding pairs of electrons around the central atom. If there are lone pairs present the geometry changes into one of the ten subtypes for a total of 15 VSEPR geometries.
Go back to the molecules drawn above and determine their 3D VSEPR shape.

Polarity
A polar covalent molecule results from an imbalance of charge. The molecule has a region with a partial positive charge and a region with a partial negative charge. A polar molecule results when bond dipoles do not counteract one another. To find out if a molecule is polar or non-polar it must be drawn in 3D first.
Ex.
Intermolecular Forces
The forces of attraction and repulsion between particles (intermolecular) determine the physical properties of a compound. The more intermolecular forces the higher the melting/boiling point.
1) Hydrogen Bonding
- strong intermolecular attraction between H from an N-H, O-H, or F-H group on one molecule and an
N, O or F on another molecule
- molecules that can form hydrogen bonds with themselves have a higher boiling point than similar molecules that cannot
- molecules that can form H-bonds with water are usually soluble in water
2) Dipole-Dipole Interactions
- attractive forces between polar molecules
- polar molecules usually have a higher boiling point than similar non-polar molecules
- polar molecules that can form H-bonds have even higher boiling points
- if a molecule contains a long non-polar section it will be less polar and less soluble in water than a similar molecule with a smaller non-polar section
3) Dispersion Forces
- relatively weak attractive forces that occur between all covalent molecules
- strengthen as size of molecule increases so a molecule with a greater number of carbon atoms usually has a higher boiling point than a molecule with fewer carbon atoms
Homework: pg. 79 #13-18, pg. 82 #9-12,14,16, read pg. 80-81