Chemistry 20 H
Chapter 1 & 2 Review
Resource:
In-class handouts
Chemistry The Central Science
Objectives:
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.
Recognize and interpret domestic hazard symbols.
Recognize and interpret WHMIS symbols.
Recognize various pieces of lab equipment.
List safety procedures.
Be able to classify matter. (1.3)
List the important base and derived units of the metric system. (1.4)
List the important prefixes in the metric system. (1.4)
Be able to convert between Celsius and Fahrenheit. (1.4)
Be able to perform density calculations. (1.4)
Use and perform calculations in scientific notation. (notes)
Compare and contrast between accuracy and precision. (1.5)
Use significant digits in calculations. (1.5)
Perform dimensional analysis. (1.6)
Discuss the nature and roles of the proton, neutron and electron. (2.3)
Discuss the size of the atom. (2.3)
Identify an element based on number of protons. (2.3)
Calculate ion charge from the number of protons and electrons. (2.3)
Calculate the mass number based on the number of protons and neutrons. (2.3)
Use atomic symbols to indicate atomic number, mass number and atomic charge. (2.3)
Calculate average atomic mass of an element based on isotopic abundance. (2.4)
Predict the properties of elements based on their position on the Periodic Table (2.5)
Compare and contrast between molecular and ionic compounds. (2.6, 2.7)
Predict ion charges based on position on the Periodic Table. (2.7)
Be able to predict the formulas of ionic compounds based on ionic charge. (2.8)
Be able to name ionic compounds, including those involving polyatomic ions and acids. (2.8)
Be able to write formulas of ionic compounds, given their systematic names. (2.8)
Be able to name binary molecular compounds (2.8)
Be able to write formulas of binary molecular compounds, given their systematic names. (2.8)
Be able to name simple organic compounds, given their structures or formulas. (2.9)
Be able to write the structures or formulas of simple organic compounds, given their names.
(2.9)
Vocabulary
acute
chronic
flammable
radioactive
base unit
corrosive
metric prefix
reactive
biohazard
derived unit
oxidizer
carcinogen
explosive
poisonous
see Summary and Key Terms, page 30 and 69-70.
First Thing !
Review metric base units, derived units and prefixes; MEMORIZE !
2
Lab Safety
Hazardous Materials in the Home

see attached handout for symbols
Two outlines:
1. 8-sided:
2. Triangle:
Poison
- skull and crossbones symbol
- poisons can enter the body in one of four ways:
a)
b)
c)
d)
- poisons can cause anything from mild illness to death, depending on the nature of the poison
Corrosive
Flammable
Explosive
Radiation
- symbolized by a hand eaten away to the bones in a beaker of liquid
- symbolized by a flaming surface
- symbolized by an exploding bomb
- symbolized by a 3-sided fan (trefoil)
-
WHMIS

Workplace Hazardous Materials Information System

WHMIS is a system of warning symbols and information sheets which detail the danger, safe handling and
disposal of a variety of chemical substances in Canada.

All chemicals handled in Canada must be labeled using the WHMIS system.

There are 8 classes of hazards, each with its own symbol. Be familiar with symbol, class (division) and
general hazard.
3
Toxicity
Acute Toxicity
-
Chronic Toxicity
-
Biohazard
-
In Case of An Accident:
Inhaled Poison
Contact of Poison with
Skin or Eyes
Swallowed Poison
Swallowed Corrosives
Lab Procedures and Rules
1.
2.
3.
2.
No eating or drinking in the lab.
Do not wear loose clothing during a lab. Tie long hair back.
Do not sit on the lab bench; you do not know how clean it is.
Treat all chemicals as if they were hazardous:
- never taste chemicals.
- wash hands after chemicals have been handled.
- wear eye protection when instructed.
5. Never perform unauthorized experiments.
6. Report all accidents immediately. Do not attempt to clean it up until checking with the teacher.
7. If you get a chemical solution in your eye, do not wait for the teacher; go to the eyewash station immediately
and wash the eye for at least 5 minutes.
8. If you get chemicals on your clothes, wash the clothes thoroughly.
9. Clean all equipment thoroughly and put it back where it belongs after a lab.
10. Clean your lab station and equipment thoroughly after a lab.
Laboratory Equipment. Be able to identify the items on the sheet given. Note: the flask is an erlenmeyer flask.
Complete lab safety quiz
4
1.2 Classifications of Matter
•
Matter is
States of Matter
•
Solids, liquids and gases are the three forms of matter called the states of matter.
•
•
Properties described on the macroscopic level:
•
gas (vapor):
•
liquid:
•
solid:
Properties described on the molecular level:
•
gas:
•
liquid:
•
solid:
Pure Substances
•
Atoms
•
Elements
•
Compounds
•
Law of Constant (Definite) Proportions (Proust):
•
Mixtures:
•
•
5
•
•
Heterogeneous mixtures do not have uniform composition, properties, and appearance, e.g., sand:
o
suspension
o
emulsion
o
colloid
solutions:
M a tte r
M ix tu re
P u re S u b s ta n c e
E le m e n t
H om ogeneous
S o lu tio n
C o llo id
C om pound
H e te ro g e n e o u s
S u s p e n s io n
E m u ls io n
1.3 Properties of Matter
•
Each substance has a unique set of physical and chemical properties.
•
Physical properties
•
Chemical properties
•
Characteristic Physical Properties
•
Properties may be categorized as intensive or extensive.
•
o
Intensive properties
o
Extensive properties
Intensive properties give an idea of the composition of a substance whereas extensive properties give an
indication of the quantity of substance present.
6
Physical and Chemical Changes
•
Physical change
•
Chemical change
Separation of Mixtures
•
Key: separation techniques exploit differences in properties of the components.
•
Filtration:
•
Distillation
•
Chromatography: exploit solubility of components.
Complete questions 1.16 to 1.22, even
1.4
System Internationale (the modern Metric System)

see the attached handout for prefixes, base and derived units.
Error in Science
Random Error
Systematic Error
Base Units
7
Derived Units
•
All equations must be dimensionally consistent
•
Ex:
•
Prefixes are used to change SI units by powers of 10.
SI Units
d = 1/2at 2
Dimensional Analysis

Convert 72.3 mg to grams

Convert 4.23 L to mL

Convert 21.1 μA to hA

Convert 55 m 2 to cm 2

Convert 6120 mm 3 to m 3

how many seconds are there in 5.00 days?

Calculate the velocity of a car in m/s if it travels at 110 km/h
•
If a car is going 50 mile/hr convert to km/hr and m/s
Practice Problems
1. 11.2 m to mm
2. 0.25 cA to µA
3. 3.1 MN to kN
4. 768 mL to daL
5. 582 cm 3 to m 3
6. 1.18 dm 2 to mm 2
7. How many seconds in a leap year?
8. Convert the speed 5.30 m/s to km/h
8
Significant Figures
• Remember all significant figure rules learned previously:
•
2 cases of infinite significant figures:
•
Count the significant figures:
1.
18.56 m
9.
15 000 000 A
2. 1500 ºC
10.
406.010 mol
3.
11.
120. mm
4. 0.0062 L
12.
0.00920 g
5.
0.0128 km
13.
2.300 kPa
6.
20 apples
14.
500 students
7. 8.0 J
15.
100 000 t
8. 1.03 x 10 4 N
16.
90 502 cm
0.5306 kg
Rounding Rules
• The procedure for dropping off digits in a number is called rounding off.
1.
2.
Example:
14.2481 kg rounded to three digits is _________
Example:
7.8361 km rounded to three digits is __________
4.25501 s rounded to three digits is __________
Example:
3.475 m rounded to three digits is __________
3.485 m rounded to three digits is __________
3.
9
Round these:
•
6.249 mm, 2 s.d.
56087250 N, 4 s.d.
•
10.98 g, 3 s.d.
21.35 m, 3 s.d.
•
0.0573 mol, 2 s.d.
450.5 kL, 3 s.d.
•
69.95 km  h-1, 2 s.d.
67.77 mg, 1 s.d.
•
298.036 cm 3, 4 s.d.
2800 L, 3 s.d.
•
349.9 A, 3 s.d.
675 J, 2 s.d.
•
9.100 g, 2 s.d.
Significant Figures in Mathematical Calculations
1. Addition and Subtraction
Examples: Calculate the following using the correct number of significant figures.
2.
3.
4. Multiplication and Division
Examples: Calculate the following using the correct number of significant figures.
10
Scientific Notation
• Used as a shorthand method of writing large or small numbers
•
Multiplication and Division of Scientific Notation

Addition and Subtraction of Scientific Notation
a) 9.25 g + 4.10 g - 2.05 g =
b) 134.8 mL + 2.05 mL - 13 mL =
c) 14.896 A - 2.42 A + 4.60 A =
d) (3.45 x 10 -1 m) - (4.789 x 10 -3 m) =
e) (7.9 x 10 -2 N) + (2.05 x 10 -1 N) =
f) 4.18 cm3 x 0.051 960 g/cm 3 =
g) 0.50 g ÷ 4.12 g/mol =
h) (9.330 x 10 -2 m) x (4.612 x 10 1 m) =
i) (1.981 x 10 1 mol) ÷ (2.5 x 10 -2 mol/L) =
j) (4.68 x 10 -4 kg) x (8.743 x 10 5 m) =
(1.04 x 10 -2 s) 2
Solving Problems
• Three Steps
1.
2.
3.
Example 1
• Current and voltage are related by the equation V = IR. A 12 V car battery is connected to a 3 
brake light. What is the current carrying energy to the lights?
1. Quantities
2. Formula
3. Manipulate 4. Substitute and Solve – check units!!
11
Example 2
The potential energy, PE, of a body of mass, m, raised to a height, h, is expressed mathematically as PE =
mgh, where g is the gravitational constant. If m is measured in kg, g in m/s 2, h in m, and PE in joules, then
what is 1 joule described in base units?
Example 3
• You are cracking a code and have discovered the following conversion factors:
• 1.23 longs = 23.0 mediums and 74.5 mediums = 645 shorts.
• How many shorts are equal to 1 long?
Example 4
• You are given a rectangular bar where the length = 2.347 m, thickness = 3.452 cm,
height = 2.31 mm, mass = 1659 g.
• determine the volume in cubic meters
• determine the density in grams per cubic centimeter and kilograms per cubic meter.
Temperature



o
o


12
Conversion from Celsius to Fahrenheit to Kelvin Temperature Scales:

a)
b)
c)

a)
b)
•
a temperature of 371 K is measured
• what is this temperature in degrees Celsius and Farenheit?
•
a temperature of -40.0 °F is measured
• what is the temperature in °C and K?
Complete questions 1.33 & 1.34
Density




Calculate the density of a substance if it has a mass of 16.36 g and a volume of 1.21 cm 3
Calculate the volume of 3.56 kg of aluminum
13
Calculate the mass of a pure gold ring with a volume of 0.762 cm 3
Complete questions 1.28, 1.30, 1.32
The Atom
A Bit of History
Law of Conservation of Mass (Antoine Lavoisier, 1743-1794)
Law of Constant Composition (Joseph Proust, 1754-1826)
John Dalton (1766-1844)
 Chemical Atomic Theory:
a)
b)
c)
d)
e)
f)
g)
J.J. Thomson (1856-1940)








14
Robert Millikan (1868-1953)

famous for the Millikan Experiment



Henri Bequerel (1852-1908)


Ernest Rutherford (1871-1937)

discovered 3 types of radiation:
o
o
o

performed the gold foil experiment; he shot alpha particles (massive and positively charged) at a gold
foil:
o
o
o

the experiment had 3 main conclusions:
a)
b)
c)
Other Subatomic Particles
Nucleus
 is the centre of the atom. It is made up of neutrons and protons.
15
Property
Proton
Location
Mobility
Charge
Mass
Role In Atom



The angstrom :
Size of nucleus:
Atomic mass unit:
Atomic Charge


An atom
Charge happens because:

To calculate ion charge:


Net atomic charge =

Nitrogen

Strontium

Se, 36 electrons

Mn, 18 electrons

53 p+, 54 e-

88 p+, 86 e-
Neutron
Electron
16

If electrons leave or are added to an atom, the atom becomes an ion:
- a positive ion is a
- a negative ion is a
Atomic Numbers, Mass Numbers, And Isotopes

Atomic number:

Mass number:

Nucleon:

Convention:

Isotopes have:

Isotopes have:

Nuclide:




Hydrogen-1 (1 proton) is called hydrogen.
Hydrogen-2 (1 proton, 1 neutron) is called deuterium.
Hydrogen-3 (1 proton, 2 neutrons) is called tritium.
Isotope symbols

Chlorine, 17 p+, 18 n0

Bismuth, 83 p+, 126 n0

Scandium, 21 p+, 24 n0

Aluminum, 13 p+, 13 n0

Chlorine, 17 p+, 18 n0
•
51 protons, 70 neutrons, 54 electrons
•
56 protons, 81 neutrons, 54 electrons
Complete the Isotope and Charge Table
17
Isotopes and atomic mass

Atomic and molecular masses can be measured with

Elements are

Average atomic mass = (atomic mass A)(proportion of A) + (atomic mass B)(proportion of B)
+ …..

Proportion of isotope = Abundance (%)
100 %
Example 1: Carbon
Example 2: Sulfur
Determine the average atomic mass for boron, oxygen, titanium and selenium
The Periodic Table

Elements are arranged in order of

Periodicity is

Rows in the periodic table are called
.

Columns in the periodic table are called
. These have similar

Some of the groups in the periodic table are given special names:

Group 1 (1A):

Group 2 (2A):

Group 16 (6A):

Group 17 (7A):

Group 18 (8A):
.
18

Nonmetallic elements, or nonmetals, are located
o Properties:

Metalloids (
o
) include:
Properties:

Metallic elements, or metals, are located
o Properties:

Transition metals

Transuranium elements

Lanthanide series

Actinide series
19
Systems of Nomenclature – the naming of chemical substances
Types of Chemical Substances



Inorganic
o elements
o molecular
o ionic
Organic
o aliphatic
o cyclic
o substituted
Common substances
o inorganic or organic
All naming is based on the same 2 principles:



Here is the list in a bit more detail:
1.
2.
Elements –

monoatomic:

diatomic:

polyatomic:
Common names. See the attached list.

3.
be familiar with:
Compounds:
-
Molecular Compounds:


Prefixes:
1
6
2
7
3
8
4
9
5
10
20
Examples:

N2O5

CO2

CO

Cl2O7

tetraphosphorus decaoxide

tetrasulfur tetranitride

bromine monofluoride

diarsenic triselenide
-ides:
Ionic Compounds







K1+
F1-

Mg2+
Cl1-

Al3+
S2-

Ca2+
S2-

Sr2+
OH1-

NH41+
SO42-
Complete Ion Charge & Formula Worksheet
Simple Ionic Compounds








21
Naming Simple Ionic Compounds



NaCl

CaF2

Al2S3

Ag3P

cadmium chloride

strontium nitride
Complex Cations (Stock System)




Formula Charge of Cation Charge of Anion
Name of Compound
SnF2
SnF4
CoF3
Manganese (IV) oxide
Uranium (VI) sulfide
Cations to Know (i.e. memorize)
22
Classical Nomenclature


o

o



This system is used the same way as the Stock system. This table will show you how the classical names compare
to the Stock names:
Formula Charge of Cation Charge of Anion Stock Name
SnF2
Sn2+
F1-
Tin (II) fluoride
SnF4
Sn4+
F1-
Tin (IV) fluoride
2+
1
Cobalt (II) chloride
1
CoCl2
Co
Cl -
CoCl3
3+
Co
Cl -
Cobalt (III) chloride
Cu2O
Cu1+
O2-
Copper (I) oxide
2+
2-
Copper (II) oxide
Br
1-
Mercury (I) bromide
Br
1-
Mercury (II) bromide
CuO
Cu
Hg2Br2
Hg22+
HgBr2
Hg
2+
O
Polyatomic Ions

Ions that contain more than 1 atom.

3 cations:

2 types of anions:


Oxyanions


o
o
o
o
Write the examples on the table:
Classical Name
23

if more than two members in a series:
o
o




o

Adding Hydrogen
o any anion with a charge of 2- or greater can add hydrogen:






Other Anions:
o
o
o
o
o
o
o
o
o
Names and Formulas of Acids
 Acids are substances that yield hydrogen ions when dissolved in water (Arrhenius definition).
 The names of acids are related to the names of anions:
o

o

o

Waters of Hydration



24
o

o

Complete the inorganic nomenclature worksheet
25
Organic Compounds
Compounds made of:
Known as
Carbon makes
3 types of organics in this class:



Naming Organic Compounds
Naming hydrocarbons involves looking at 3 aspects:
1.
2.
3.
There is a system for naming chains of carbon compounds. It consists of two steps:
-
a prefix tells you how many carbons are in the chain.
a suffix gives you information about that carbon chain.
Length of Carbon Chain
• how many carbons are bonded in a chain gives the prefix of the name:
1
2
3
4
5
6
methethpropbutpenthex-
7
8
9
10
11
12
heptoctnondecundecdodec-
Other prefixes include:
Number of Carbon Atoms
11
12
13
14
15
16
20
25
30
40
50
100
Prefix
undecdodectridectetradecpentadechexadeceicospentacostriaconttetracontpentaconthect-
26
Family Background
There are three main groups of hydrocarbons, based on the type of bonds which join the carbons:
Alkanes
-
Alkenes
-
Alkynes
-
27
Isomers
•
•
structure associated with a chemical formula.
it is possible to have more than one isomer for a given formula:
•
•
•
Location of Multiple Bond
• number the carbon chain from the end closest to the double or triple bond; can be left-to-right or
right-to-left:
Functional Groups
•



1. Hydrocarbon groups
•
Hydrocarbon chains are carbon chains attached to the main chain:
•
Prefixes are used to indicate the number of carbons in the chain:
28
•
numbering is used when isomers are possible:
•
numbering for alkenes or alkynes takes precedence over functional groups:
•
two or more functional groups are all recognized in the name:
29
2. Halogens
-
halogens (elements from group 17) can replace hydrogens on the carbon chain. Following is a list of prefixes:
Halogen Prefix
Fluorine
fluoro-
Chlorine
chloro-
Bromine
bromo-
Iodine
Iodo-
3. Alcohol Groups
•
are formed when –OH groups are attached to the main chain:
30
Cyclic and Aromatic Hydrocarbons
•
straight line hydrocarbons are called
•
carbons joined in a ring are called
•
some cyclics are called
•
the smallest ring contains 3 carbons;
•
aromatic hydrocarbons:
o
31
Other functional groups
4. Ethers (R–O–R’)
Two naming systems are possible:
1. The shorter carbon chain is the prefix; -oxy- is added in the middle and the longer carbon chain is the root
compound.
2. The shorter chain is the first prefix, the longer chain is the second prefix, followed by ether.
5.

Compounds with a Carbonyl Group
The carbonyl group is C=O.
a) Aldehydes


have the oxygen double bonded to a carbon at the end of the chain.
the name is the root chain, followed by suffix -al
b) Ketones


have the oxygen double bonded to a carbon in the middle of the chain.
the name is the root chain, followed by suffix –one. Numbering is used, if necessary.
32
c) Carboxylic Acids


d)

The carboxyl functional group is –COOH.
the name is the root chain, followed by suffix –oic acid.
Esters are a combination of an organic acid and an alcohol.
Esters are named using the alcohol part first and then the acid part.
 Example: The ester formed from ethanol and acetic acid is ethyl acetate (also called ethyl
ethanoate.
 sometimes the acid part is given the ending -oate
33
e)
Amines and Amides


Amines are organic bases.

Amides look like an ester, but with a nitrogen atom in carbon chain instead of the oxygen:
they include a nitrogen in the carbon chain
34
In Summary:
35
•
Geometric Isomers are differences in shape resulting from arrangements around a double bond on an
alkene or in a cyclic compound.
o

the arrangements are designated by the prefix cis- or trans-
optical isomers are more subtle arrangements, allowing compounds to rotate light in different ways.
o rotation is either left (levo-) or right (dextro-)
Complete organic nomenclature assignment