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IB CHEMISTRY YEAR 1 SUMMER ASSIGNMENT

The IB Chemistry Year 1 Summer Assignment is a review of chemistry concepts covered

in Grade 9 CPF (Chemistry & Physics Fundamentals) course at ASD. It will take

between 4-5 hours for students who have taken CPF before. It will take longer for

students who have no Chemistry background.

I. Read class notes & complete sample problems (1.5 hours)

II. Complete IB Chemistry Y1 Summer Assignment. (2 hours)

III. Correct assignment with IB Chemistry Y1 Summer Assignment Answers (0.5

hours) and come to class with questions!

Completion of assignment will be checked on the first day of school. There will be ONE

review class. If students are unable to understand these concepts on their own, they will

be required to attend sessions outside of class time during the first week of school (Tues,

Aug 27 and Wed, Aug 28 after school). A test will be given on Sun, Sept 1 or Mon, Sept

2 which will have a similar format to the assignment.

All materials are available on pdf on www.asd.edu.qa

Email any questions to jlamontagne@asd.edu.qa

Class notes (with links)

Units & Measurement/ Significant Figures and Calculations

A. Scientific Notation

B. Dimensional Analysis

C. SI Units- The Metric System & Metric Conversion

D. Significant Figures

E. Significant Figures & Calculations

Molecules & Ions / The Modern View of Atomic Structure

F. Classification of Matter

G. Definitions to Classify Matter

H. Physical & Chemical Properties & Changes

I. The Modern View of Atomic Structure

An Introduction to the Periodic Table

J. The Modern Periodic Table

K. Parts of the Periodic Table

L. Using the Periodic Table

M. Chemical Symbols

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Naming Simple Compounds

N. Introduction to Naming Compounds

O. Naming Binary Ionic Compounds

P. Writing Formulas for Binary Ionic Compounds

Q. Naming Ionic Compounds with Transition Metals

R. Writing Formulas for Ionic Compounds with Transition Metals

S. Naming Ionic Compounds with Polyatomic Ions

T. Writing Formulas for Ionic Compounds with Polyatomic Ions

U. Naming Binary Covalent Compounds

V. Writing Formulas for Binary Covalent Compounds

W. Chemical Nomenclature Flow Chart

Balancing Chemical Equations

X. Introduction to Chemical Equations

Y. Rules for Writing Chemical Equations

Z. Rules for Writing Balanced Chemical Equations

AA. Types of Chemical Equations

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USEFUL WEBSITES

1. Scientific Notation: http://janus.astro.umd.edu/cgi-bin/astro/scinote.pl

2. Significant Figures- http://janus.astro.umd.edu/cgi-bin/astro/scinote.pl;

http://science.widener.edu/svb/tutorial/sigfigures.html

3. Atomic Structure http://science.widener.edu/svb/tutorial/protons.html

4. The Periodic Table - http://www.learner.org/interactives/periodic/index.html

(information and interactive problems with answers)

5. Balancing Chemical Equations http://education.jlab.org/elementbalancing/index.html

6. On line Balancer of Chemical Equations http://www.webqc.org/balance.php

7. For videos on almost anything go to http://www.khanacademy.org/

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A. SCIENTIFIC NOTATION

1. Scientific notation is a method of expressing large or small numbers as factors of the power

of 10.

It is more compact

Easier to understand

Easier to manipulate

Ex: 0.0000000013 m = 1.3 x 10-9

m or 7500000 g = 7.5 x 106 g

2. The form for scientific notation is:

a x 10b

a = real number between 1- 9.99…. (but not equal to 10)

b = positive or negative integer

3. For large numbers, exponents are positive. For example for 7500000:

a = 7.5 (a number between 1-10)

To find b, count the number of spaces to the right of the decimal point. There are 6

places to the right (+6).

7500000= 7.5 x 106

4. For small numbers, exponents are negative. For example for 0.0000000013:

a =1.3 (a number between 1-9.99….)

To find b, count the number of spaces to the left of the decimal point. There are 9 places

to the left (-9).

0.0000000013= 1.3 x 10-9

5. Multiplying numbers with scientific notation:

Exponents are added

Ex: 5.00x 10-8

x 6.02 x 1023

= (5.00 x 6.02) x 10(-8+ 23)

= 30.1 x 1015

= 3.01 x 1016

6. Dividing numbers with scientific notation:

Exponents are subtracted

1.76 x 104 = (1.76) x 10 (

4-23) = 0.292 x 10

-19 = 2.92 x 10

-20

6.02 x 1023

6.02

Please note that many calculators and computers cannot conveniently use this notation so

the E notation is used ex: 2.92 E-20 instead of 2.92 x 10-20

. It is expected that all answers

on a quiz/ test are written in a x 10b

format.

For practice, go to: http://janus.astro.umd.edu/cgi-bin/astro/scinote.pl

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B. DIMENSIONAL ANALYSIS = FACTOR LABEL METHOD

1. All measurements must have units. The measurement 12.8 does not exist. But the measurement

12.8 m, 12.8 mL, or 12.8 cg do exist.

2. Units involved in calculations follow all the rules of algebra. In algebra, the following

simplification would be performed.

x2 • y = x • y

x

Similarly, m2 • g = m • g

m

3. All calculations involving measurements must have units. The only exception is when a ratio of

identical units is simplified. For example:

28 g = 4

7 g

4. All conversion factors must have units. Conversion factors are always expressed as the quotient

of two units. For example, to convert 5 minutes to seconds, the conversion factor 60 s/1 min is

used.

5 min x 60 s = 300s

1 min

Sample Problem 1 Determine the units of the answer to the following problem:

12 m x 100 cm = 1200 ?

1 m

Sample Problem 2 Determine the units of the answer to the following problem:

55.0 g = 5.5 ?

5.0 cm x 2.0 cm x 1.0 cm

5280 m x 1 km = 5.28 ?

1000 m

12.5 cm x 10.0 cm = 125 ?

5. Many of the problems in chemistry require the conversion from a given unit to another unit. The

simplest way to attack this type of problem is to use the appropriate conversion factor.

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6. A conversion factor is a fraction formed from a valid relationship and is used to

switch from one unit to another.

(given quantity) x (conversion factor) = (desired quantity)

Example: 1 stone = 14 lbs (pounds) and 1 kg = 2.2 lbs. Express 48 kilograms in stones.

7. An advantage to the factor label method is that it actually lets you know if you have

done the wrong arithmetic by looking at the units!

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C. SI UNITS - THE METRIC SYSTEM & METRIC CONVERSION

1. A standard system of units is essential if measurements are to be made consistently.

2. The SI (French for International System of Units) or metric system is the dominant

system of units in science and engineering.

SI BASE UNITS

MEASUREMENT UNIT SYMBOL

Length

Mass

Time

Temperature

Amount of substance

3. When making measurements, we sometimes find that the basic unit is too small or too

large. Using the multipliers eliminates the need for scientific notation.

4. Below is a metric conversion line to help make conversions in metrics quick and

easy.

5. The basic unit is in the center (ex: meters, liters, grams, seconds, etc.)

6. When converting from one metric value to another, count the number of units (or the

power) you move either up or down the line and then move the decimal point the

same number of spaces. The arrow indicates the direction you move the decimal

point. For squared unit, multiply number of spaced by 2. For cubed units, multiply

number of spaces by 3. Ex: 1 cm2 = ________mL

2 or 1 m

3 = ________________ cm

3

DECIMAL MULTIPLIERS

Prefix Giga Mega kilo hecto deka base deci centi milli micro

Symbol G M k h da d c m µ

Multiplier 109 106 103 102 101 10-1 10-2 10-3 10-6

Mnemonic to help with the order of commonly used multipliers:

Great - - Mighty - - King Henry Died By Drinking Chocolate Malted - - Milk

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7. Convert each of the following measurements into the units identified. Remember

significant figures!

a. 4.6 km = ___________________ m

b. 15.0 mL = __________________ daL

c. 98.4 mg = ___________________ kg

d. 0.008760 ML = _____________________ L

e. 3.45 kg = ___________________ g

f. 42 m3 = ______________________ cm

3

The following conversion factors will be very useful in making metric conversions.

1000 mm or 1 m

1 m 1000 mm

100 cm or 1m

1 m 100 cm

1000 m or 1 km

1 km 1000 m

These conversion factors can be used for grams and liters as well as meters.

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D. SIGNIFICANT FIGURES

1. During laboratory sessions you will be required to measure certain quantities such as

mass, length, temperature or volume. How we treat those numbers is important as a

scientist does not want to imply any more or any less precision in experimental results

than is actually there.

2. All measurements have a certain degree of uncertainty associated with them. When

scientists report measurements they report all the digits that are known plus one

uncertain digit.

3. All the digits in a measurement expression that are known with certainty (the smallest

division), plus the first digit that is uncertain (estimated) are called significant

figures.

4. Significant figures indicate the uncertainty of a measurement. The measurement 4.72

cm is precise to the second decimal place. The digit 2 is the last significant figure and

the first uncertain digit. 4.72 cm contains three significant figures.

5. Exact numbers (i.e. 12 atoms) are said to have an infinite number of significant

figures.

6. All nonzero digits in a measurement are always significant.

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7. Zero, however, is not always a digit. Sometimes, zero is a placeholder. When a zero is

a placeholder, it is not a significant figure.

8. The following rules will assist you in determining whether a zero is a significant

figure or a placeholder.

RULE MEASUREMENT

EXPRESSION

SIGNIFICANT

FIGURES

1. All nonzero digits are significant. 83.591 m

5

2. All zeros between two nonzero digits are significant. 5007 L

10.0005 g

136.05

4

6

5

3. Trailing Zeroes with no decimal point: Zeros to the

right of a nonzero digit, with no decimal point are not

significant.

200,800 km

2100 L

1,000,000 g

4

2

1

4. Trailing Zeroes with a decimal point: All zeros to the

right of a decimal point and to the right of a nonzero digit

are significant.

40.00 g

0.005070 kg

40.

4

4

2

5. Leading Zeroes: All zeros to the right of a decimal point

but to the left of a nonzero digit are not significant. A lone

zero to the left of a decimal point is never significant.

0.00012 g

0.853 m

2

3

Sample Problems:

a. How many significant figures are there in 21.589 m?

b. How many significant figures are there in 28005 km?

c. How many significant figures are there in 0.00025 kg?

d. How many significant figures are there in 23,000 L?

e. How many significant figures are there in 80.0 cm?

Sample Problem Answers: 5; 5; 2; 2; 3

Practice Problems go to http://science.widener.edu/svb/tutorial/sigfigures.html

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E. SIGNIFICANT FIGURES & CALCULATIONS

Suppose you wished to multiply 24 cm by 318 cm. How many significant figures should

the answer contain? The result of calculations involving measurements can only be as

precise as the least precise measurement. In the above problem the answer can only have

two significant figures. The following rules will enable you to determine the number of

significant figures in the result of calculations involving measurements.

1. Rule 1- Multiplication and Division: The answer contains the same number of

significant figures as the measurement with the least number of significant figures.

NOTE: The position of the decimal point does not determine the precision of the

answer.

Sample Problem 1

Determine the precision of the product of 24 cm x 31.8 cm.

Sample Problem 2

Determine the correct number of significant figures for the quotient of 8.40 g ÷ 4.2 mL.

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2. Rule 2- Addition and Subtraction: The sum or difference has the same number of

decimal places as the measurement with the least number of decimal places.

NOTE: The position of the decimal point determines the precision of the answer.

Sample Problem 3

Determine the precision of the sum of 49.1 g + 8.001 g.

Sample Problem 4

Determine the precision of the difference of 81.350 m - 7.35 m.

3. Rule 3- When there are a series of calculations to obtain a final result, DO NOT

ROUND OFF UNTIL THE END. Or if you do round off, leave at least one extra

digit until the end of all the calculations. For rounding:

If the number is less than 5, then the preceding digit stays the same ex: 1.34 to

1.3 for 2 sf

If the number is equal or greater than 5, then the preceding number is

increased by 1. ex: 1.35 to 1.4 for 2 sf

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F. CLASSIFICATION OF MATTER

1. Matter: is defined as anything which occupies space and has mass.

2. All types of matter can be classified under the three headings: solids, liquids and

gases.

These three headings are known as the three states of matter.

Note: light and energy are not forms of matter.

3. Matter is made up of atoms.

4. Table: Properties of Matter

Properties Solid Liquid Gas

Arrangement between atoms

(close/ random/ far apart)

close random Far apart

Attraction

between atoms (high/ moderate/

low)

high moderate low

Movement of atoms (high/

moderate/ low)

low moderate high

Volume (fixed/ changes)

fixed fixed changes

Shape (fixed/ changes)

fixed changes changes

Compressibility (high/ low)

low low high

Diagram (of atoms)

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G. DEFINITIONS TO CLASSIFY MATTER

1. Atoms: matter is built up from very tiny units called atoms.

There are 118 different types of atoms (elements)

All matter is made up of just these atoms.

Atoms vary in size and mass.

2. Pure substance: a homogeneous material consisting of one kind of substance.

Elements and compounds are pure substances.

3. Element: a pure substance that cannot be decomposed into anything simpler by

chemical means.

They contain just one type of atom.

An element is represented by a symbol and can be in any phase.

There are 118 elements; 25 do not occur in nature; 7 are diatomic (two

identical atoms bonded together).

Elements are represented by symbols (1 or 2 letters) found on the periodic

table. Ex: O: oxygen, He: helium, Na: sodium

4. Compound: a pure substance made of two or more different atoms that are

chemically combined.

There are two types of compounds: ionic compounds and molecular compounds

Ionic compound

Composed of ions (usually a metal and non-metal) because there has been a

transfer of electrons Ex: NaCl, CaCO3

Ionic compounds form large crystal lattice structures and the chemical

formula represents the simplest ratio.

Crystal Lattice of NaCl (pink = Na; red = Cl)

Molecular Compound

Composed of non-metals that share electrons

Molecular compounds form molecules that are identical to each other Ex:

CO2, H2O

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5. Molecule: made of two or atoms that are chemically bonded by a covalent bond.

The atoms can be the same (diatomic element) or different (molecular compound)

6. Mixtures: are composed of two or more substances (elements or compounds). Each

substance retains its own characteristic properties.

The substances are not chemically combined.

They can be separated by physical means (filtration, evaporation, etc.)

The substances of a mixture join in variable amounts.

Mixtures can exist in one or more phases (mud and water, oil and water)

7. Homogeneous mixture: substances are composed of a single phase and are uniform

in composition and properties throughout a given sample.

8. Heterogeneous mixture: substances are composed of more than one phase- the

components separate into different regions. Thus the composition and physical

properties vary from one part of the mixture to another.

9. Overview

Matter

Mixture

Heterogeneous Mixture

Homogeneous Mixture

(solutions)

Pure Substance

Element

Atoms

Compounds

Molecular Compound

Molecule

Ionic Compound

Crystal Lattice

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H. PHYSICAL & CHEMICAL PROPERTIES AND CHANGES

1. Physical property: can be determined without causing a change in the composition

of a substance. Examples:

State

Density

Boiling point (b.p.)

Melting point (m.p.)

Color

Hardness

Volatility (ease to boil)

2. Chemical property: the behavior of a substance during a chemical reaction (i.e.

while changing a substance to another with different composition, structure and

properties).

Flammability

Reactivity

3. Physical change: occurs when there is a change in state but not in chemical

composition. The following must hold true:

The original material is retained.

Change is reversible. Means of physical separation may be used: magnetism,

sedimentation, flotation etc.

Examples: melting, freezing, condensation, dissolving, boiling, mixing, cutting,

tearing, chopping.

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4. Chemical change: occurs when there is a change in chemical composition. The

following occurs:

Formation of a new material.

Change is irreversible by using means of physical separation.

There is usually some evidence such as a color change, change in temperature or

mass, evolution of gas or formation of a precipitate (insoluble solid).

Chemical changes can be represented using chemical equations: O2 + C → CO2

Examples: burning, cooking, digesting, rusting, corroding, growing, tarnishing.

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I. THE MODERN VIEW OF ATOMIC STRUCTURE

1. An atom is made up of a tiny dense nucleus (diameter 10-13

m) and an electron cloud

(10-8

m). This is like a pea (nucleus) in a football field (the electron cloud). Hence

the atom is empty space.

2. The nucleus contains the protons and neutrons. The electron cloud contains the

electrons (in electron shells).

Particle Symbol Charge Relative mass

(a.m.u.)

Mass

Electron e 1-

~ 5 x 10

-4 9.11 x 10

-31 kg

Proton p 1+ 1 1.67 x 10-27

kg

Neutron n 0 1 1.67 x 10-27

kg

3. The nucleus makes up for most of the atom’s mass (> 99.9%) since it contains the

protons and neutrons (electrons account for little of the atom’s mass).

4. The arrangement of the electrons in an atom is responsible for its chemical properties

since the electrons make up most of the atom’s space and thus are the part that

interacts with other atoms.

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J. THE MODERN PERIODIC TABLE

1. The number shown above each symbol is the atomic number, Z, the number of

protons in the atom.

2. The number below each symbol is the mass number, A, the number of protons and

neutrons in the atom.

3. Each element is represented by a chemical symbol, A

ZX. The atomic number is

written as a subscript and the mass number as a superscript. Ex: 12

6C for carbon,

which has 6 protons and 6 neutrons. * Note in the IB Periodic Table, it is written in

the opposite format.

4. The fundamental difference between atoms of different elements is the number of

protons in the nucleus.

5. Neutron number: to calculate the number of neutrons, subtract the atomic number

from the mass number. Ex: Na: 23-11 = 12 neutrons. For lighter elements, the

number of protons and neutrons are approximately equal. But elements with many

protons require a higher proportion of neutrons because of greater repulsion between

the large number of protons. Ex: 207

82Pb neutrons = 125

6. Isotopes: have the same chemical properties but different physical properties.

Isotopes are the same element (same atomic number) but have different mass

numbers (or different numbers of neutrons). Ex: 35

Cl and 37

Cl are isotopes. All

elements have isotopes. Natural chlorine contains 75% 35

Cl and 25% 37

Cl. This is

used to calculate the relative atomic mass which is the number recorded in the

periodic table:

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If an element is only composed of two major isotopes and the molar mass is known,

the natural abundances of the two isotopes can be calculated. Iridium (Ir) has two

isotopes 191Ir and 193Ir and has a molar mass of 192.2 g/mol. Calculate the

abundances of its isotopes.

7. Neutral atom: the atomic number is equal to the electron number. Ex: Ca has 20

protons and 20 electrons and zero charge.

8. Ion: charged atom that has either gained electrons (negatively charged) or lost

electrons (positively charged). The atomic number is not equal to the electron

number. This is indicated as a superscript on the right. Ex: Na1+

: has 11 protons, 10

electrons and a 1+

charge. Br1-

has 35 protons, 36 electrons and 1- charge.

9. Cation: a positively charged ion. i.e. Na1+

10. Anion: a negatively charged ion. i.e. Cl1-

Fill in the following table.

Symbol Atomic

Number

Mass

Number

Charge Electrons Neutrons Protons

65Cu

2+

81Br

1-

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K. PARTS OF THE PERIODIC TABLE

1. Metals are to the left of the zig-zag line (between B and Al) and non-metals are to

the right of the zig-zag line.

2. Elements along the zig-zag line are called metalloids. Ex: B, Si, Ge, As, Sb, Te, Po

3. The rows correspond to periods. The columns correspond to groups or families.

4. There are four groups that have specific family names:

Group 1: Alkali Metals

Group 2: Alkaline Earth Metals

Group 17: Halogens

Group 18: Noble Gases

5. There are 7 periods (rows). Period 6 includes the Lanthanides, and period 7 includes

the Actinides.

6. There are 18 groups (columns) which all end with a Noble Gas. The IB Periodic

Table only labels the main groups 1-8 (skips the transition metals).

7. The transition metals include elements in group 3 to12 (smaller rectangle).

8. Properties change in an orderly progression from left to right, and the pattern repeats

for each period (row).

9. Periodic law: is when chemical and physical properties of an element repeat in a

regular pattern when they are arranged in order of increasing atomic number (proton

number).

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L. USING THE PERIODIC TABLE

1. The period number corresponds to the number of energy levels for a specific

element. Ex: Sodium (Na) is in the 3rd

period, so therefore has 3 energy levels.

2. Elements in the same period have the same number of energy levels. Ex: All

elements in period 3 have 3 energy levels.

3. The group number corresponds to the number of electrons in the outer energy level

(or valence electrons). Ex: Sodium is in group 1, so it has 1 outer electron, oxygen is

in group 6, so it has 6 outer electrons.

4. Elements within the same group have similar chemical properties. This is because

they all have the same number of valence electrons Ex: Li: 1 Na:1

5. The valence electrons determine the element's chemical properties (since electrons are

the particles interacting with other atoms).

6. Metals lose electrons in chemical reactions.

7. Non-metals gain electrons in chemical reactions.

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M. CHEMICAL SYMBOLS

1. Substances are made from elements in the periodic table.

2. Each element has a unique chemical symbol containing one or two letters.

3. The first letter is ALWAYS capitalized (ex: H, He, Ca) and the second letter is

always lowercase.

4. The name of many elements come from Latin or Greek origin

Problems

1. Write the chemical symbol for the following elements. Refer to the periodic table.

a. Lead _________

b. Calcium _________

c. Sodium _________

d. Chlorine _________

e. Fluorine _________

f. Bromine _________

g. Copper _________

h. Potassium _________

i. Hydrogen _________

j. Magnesium _________

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N. INTRODUCTION TO NAMING COMPOUNDS

1. When chemistry was an infant science, there was no system for naming compounds.

Common names were used to describe compounds such as Epsom salts, milk of

magnesia and laughing gas.

2. There are nearly 5 million compounds currently known and memorizing them would

be impossible. The IUPAC (International Union of Pure and Applied Chemistry)

system is used for naming compounds.

3. There are several different types of compounds and each have different rules for

naming.

i. Simple binary (two) ionic compounds (metal + non-metal)

ii. Binary ionic compounds with transition elements

iii. Ionic compounds with polyatomic ions

iv. Inorganic covalent compounds (non-metals)

v. Organic covalent compounds (you will learn how to name these in

Organic Chemistry)

vi. Acids

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O. NAMING BINARY IONIC COMPOUNDS

1. When a metal (or positive ion) and a non-metal (or negative ion) combine together,

they form an ionic compound.

2. RULES:

i. Put the metal element (or positive ion) name first and then the non-metal

element (or negative ion) name second. i.e.: sodium chloride, calcium

fluoride, sodium hydroxide

ii. Write the metallic name as it appears on the periodic table.

iii. For the non-metallic element, take off the ending and add "IDE"

a. Nitrogen _______________

b. Oxygen _______________

c. Sulfur _______________

d. Fluorine _______________

e. Chlorine _______________

f. Bromine _______________

g. Iodine _______________

Name the following compounds.

1. LiBr2 _______________________________________________

2. Mg3P2 _______________________________________________

3. Al2S3 _______________________________________________

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P. WRITING FORMULAS FOR BINARY IONIC COMPOUNDS

1. In an ionic compound the formula represents the simplest ratio, also known as the

formula unit.

2. The oxidation number is the charge of the atom when it gains or loses electrons.

3. Rules for writing formulas for binary ionic compounds: Ex: magnesium chloride

a. Write the symbol for each element in the compound. Ex: MgCl

b. Find and write the oxidation number for each element. Refer to Table. Ex:

Mg2+

Cl1-

c. Reduce the oxidation numbers to their lowest possible ratio. Ex: Mg

2+ Cl

1-

d. Criss-cross the oxidation numbers (do not include charges) and write them as

subscripts. Ex: Mg1Cl2

e. If the subscript is a “1”, there is no need to write it. Ex: MgCl2

General Rule:

For elements in group 1 to 3: Oxidation number is equal to the group number.

For elements 5 to 7: Oxidation number is equal to (group number) - 8.

Table: Oxidation numbers of representative elements.

Group Number Oxidation Number Examples

Metals Equal to last digit

I 1 Li1+

II 2 Ba2+

III (13) 3 Al3+

Nonmetals Equal last digit - 8

V (15) -3 N3-

VI (16) -2 O2-

VII (17) -1 Cl1-

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Problems: Write the formulas of the following ionic compounds.

1. Sodium iodide

2. Rubidium oxide

3. Potassium nitride

4. Calcium sulfide

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Q. NAMING IONIC COMPOUNDS WITH TRANSITION METALS

1. Transition elements are found in groups 3-12.

2. Unlike group 1, 2, 13, 15, 16, and 17, you are unable to determine the oxidation

number of transition elements by looking at their group number.

3. Transition elements can have many oxidation numbers i.e. Fe2+

and Fe3+

. Except

silver (1+) and zinc (2

+). The oxidation number of these two elements must be

memorized.

4. Rules for writing formulas for compounds with transition elements:

a. Write the name of the transition metal. It is the same name found on the

periodic table.

b. In parentheses, write the oxidation number in roman numerals. This can be

determined by the formula.

c. For the negatively charged ion, change the ending to “ide”.

Roman numerals

Roman Numeral Corresponding Number

I 1

II 2

III 3

IV 4

V 5

VI 6

VII 7

VIII 8

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R. WRITING FORMULAS FOR IONIC COMPOUNDS WITH TRANSITION

METALS

1. Rules for writing formulas for ionic compounds: Ex: copper (II) chloride

a. Write the symbol for each element in the compound. Ex: CuCl

b. Find and write the oxidation number for each element. For the transition element

oxidation number refer to the Roman numeral in parentheses. Ex: Cu2+

Cl1-

c. Reduce the oxidation numbers to their lowest possible ratio. Ex: Cu2+

Cl1-

d. Criss-cross the oxidation numbers (do not include charges) and write them as

subscripts. Ex: Cu1Cl2

e. If the subscript is a “1”, there is no need to write it. Ex: CuCl2

Problems: Write the formula for the following compounds.

1. Copper (II) sulfide

2. Cobalt (III) iodide

3. Zinc oxide

4. Manganese (VII) fluoride

5. Chromium (VI) iodide

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S. NAMING IONIC COMPOUNDS WITH POLYATOMIC IONS

1. Some elements commonly exist bonded together in what is called a polyatomic ion

(or radical).

2. A polyatomic ion is an ion that has two or more different elements. Polyatomic ions

are treated as elements. The formula always stays the same (like a symbol). Ex: NO3

cannot change to NO2

3. For example in calcium sulphate, the sulphate part is not found as an element in the

periodic table, it is a polyatomic ion and it must be memorized. The formula is SO4

and it has a oxidation number of 2-.

4. Rules for naming ionic compounds with polyatomic ions:

a. If the element is from the periodic table follow the rules for simple ionic

compounds (first element has the same name, the second element name

changes to “ide” ending). Ex: NaOH = sodium hydroxide or NH4Cl =

ammonium chloride b. If it is a polyatomic ion, simply write the name. These names and formulas

have to be memorized!!

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Table: Polyatomic Ions

Name of Ion Formula

1. acetate C2H3O21- or CH3COO1-

2. ammonium NH41+

3. carbonate CO32-

4. hydrogen carbonate HCO31-

5. chlorate ClO31-

6. chlorite ClO21-

7. chromate CrO42-

8. dichromate Cr2O72-

9. cyanide CN1-

10. hydroxide OH1-

11. hypochlorite ClO1-

12. nitrate NO31-

13. nitrite NO21-

14. perchlorate ClO41-

15. permanganate MnO41-

16. phosphate PO43-

17. hydrogen phosphate HPO42-

18. dihydrogen phosphate H2PO41-

19. sulfate SO42-

20. hydrogen sulfate HSO41-

21. hydrogen sulfide HS1-

22. sulfite SO32-

23. hydrogen sulfite HSO31-

24. thiocyanate SCN1-

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Problems: Write the name for the following compounds.

1. NaOH ______________________________

2. K3PO4 ______________________________

3. Ba(ClO4)2 ______________________________

4. NH4F ______________________________

5. Cs2CO3 ______________________________

6. CaSO4 ______________________________

7. HCN ______________________________

8. Fe(NO3)3 ______________________________

9. Fe(NO2)2 _____________________________

T. WRITING FORMULAS FOR IONIC COMPOUNDS WITH POLYATOMIC

IONS

Rules for writing formulas for ionic compounds: Ex: Ammonium carbonate

1. Write the symbol for each element/ polyatomic ion in the compound. Ex: NH4CO3

2. Write the oxidation number for each element/ polyatomic ion. These must be

memorized. Ex: NH41+

CO32-

3. Reduce the oxidation numbers to their lowest possible ratio. Ex: NH41+

CO32-

4. Put the polyatomic ions in brackets. (NH41+

)(CO32-

)

5. Criss-cross the oxidation numbers (do not include charges) and write them as

subscripts outside the brackets. Ex: (NH4)2(CO3)1

6. If the subscript is a “1”, there is no need to write it the one or to put brackets around

the polyatomic ion. Ex: (NH4)2CO3

Problems: Write the formulas for the following

1. ammonium iodide

2. Zinc sulfite

3. Iron (II) carbonate

4. Cobalt (III) nitrate

5. Nickel (II) perchlorate

6. Manganese (IV) cyanide

7. Copper (I) phosphate

8. Chromium (II) oxide

9. Tin (IV) sulfide

10. Cerium (III) hydroxide

U. NAMING BINARY COVALENT COMPOUNDS

1. Covalent compounds occur between two non-metals. i.e. CO2

2. Since different kinds of molecules can exist between the same elements (ex: CO,

CO2) a different naming system is used.

Rules for naming binary covalent compounds:

3. Write the name of the first element. It is the same name found on the periodic table.

To determine which element goes first:

The first element is always the one furthest to the left (ex: N is written before O)

If the elements are in the same group, the first element is always further down the

group (ex: S is written before O).

4. For the second element, change the ending to "ide".

5. Add a prefix to the name of each element to indicate the number of atoms for each

element.

6. The prefix "mono" is not used for naming the first element (ex: carbon dioxide, not

monocarbon dioxide).

7. When the vowel combinations o-o and o-a appear next to each other, the first letter is

omitted to simplify pronunciation (ex: nitrogen monoxide, not nitrogen monoxide).

Table: Prefixes.

PREFIX NUMBER OF ATOMS

Mono- 1

Di- 2

Tri- 3

Tetra- 4

Penta- 5

Hexa- 6

Hepta- 7

Octa- 8

Nona- 9

Deca- 10

Problems: Name the following compounds.

1. CO

2. CO2

3. CCl4

4. IF7

5. N2O

6. SO2

7. CH4

8. NH3

9. N2O4

10. H2O

V. WRITING FORMULAS FOR BINARY COVALENT COMPOUNDS

Rules for writing formulas for binary covalent compounds: i.e.: silicon tetrafluoride

a. Write the symbol for each element in the compound. i.e.: SiF

b. Add the appropriate subscript that corresponds to the prefix name. i.e.: SiF4

Problems: Write the formulas of the following ionic compounds.

1. Disulfur dichloride

2. Carbon disulfide

3. Sulfur trioxide

4. Tetraphosphorus decoxide

5. Carbon tetrachloride

6. Hydrogen fluoride

7. Fluorine

8. Nitrogen monoxide

9. dinitrogen monoxide

10. iodine pentafluoride

W. CHEMICAL NOMENCLATURE FLOW CHART

X. INTRODUCTION TO CHEMICAL EQUATIONS

Chemical equations illustrate what is made (products) when certain chemicals (reactants)

are combined. Like a cooking recipe, where certain amounts of ingredients are required

to produce a set amount of food, chemical reactions require a certain amount of reactants

to get a desired quantity of product. This relationship is seen through a chemical

equation.

A chemical equation represents chemical reactions, which may be indicated by:

Color change

Precipitate formed (solid forming in solution)

Odor produced

Gas released

2Al (s) + 3CuCl2 (aq) 2AlCl3 (aq) + 3Cu (s)

There are several parts to a chemical equation:

1. Reactants: Substances on the left-hand side of the equation.

2. Products: Substances on the right-hand side of the equation.

3. Yield sign: → means “produces”

4. State symbols: letters indicating the state of the substances. (s) = solid; (l) = liquid;

(g) = gas; (aq) = aqueous or dissolved in water or a solution

5. Subscripts: numbers to the lower right of a chemical symbol, used to indicate the

number of atoms for that element i.e.: H2O. A subscript outside a bracket multiplies

all the elements in the bracket. i.e. Ca(NO3)2

6. Coefficients: numbers to the left of a chemical substance, they indicate the number of

molecules ex: 3Mg(OH)2.

7. The Law of Conservation of Mass: Matter can neither be created nor destroyed. In

other words the number of each type of atom on the reactant side should be equal to

the product side.

Exercise #1: From the following equation, give an example of a reactant, product,

subscript, and coefficient.

a. 2H2 (g) + O2 (g) 2H2O (l)

b. H2SO4 (aq) + 2 NaOH (aq) Na2SO4(aq) + 2 H2O (l)

Exercise #2: Indicate how many atoms of each element are in each of the following:

Na2CO3

K2CrO4

Ca3(PO4)2

3NaCl

2Ca(NO3)2

3Fe2(SO4)3

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Y. RULES FOR WRITING WORD EQUATIONS

1. Write the names of the reactants and add a + sign between reactants.

2. Draw an arrow.

3. Then write the names of the products and add a + sign between the products.

4. Add letter(s) in parentheses to indicate each substance’s state.

Exercise #1: Write the word equation for the following.

An aqueous solution of barium chloride is added to a sodium nitrate solution to produce a barium

nitrate solution and a sodium chloride solution.

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Z. RULES FOR WRITING BALANCED CHEMICAL EQUATIONS

A. Write the chemical symbols/ formulas for each substance.

a. For elements write the symbol on the Periodic Table.

b. For elements check to see if they are diatomic: Friday Night I Hang Out at the

British Club- F2, N2, I2, H2, O2, Br2, Cl2. (Note: S8 and P4 also exist in molecular form

although many times they are just written as S and P respectively.)

c. For compounds, find the correct formula by:

Finding the oxidation numbers.

If possible, simplify ratio i.e. 2:2 is reduced to 1:1 or 4:2 is reduced to 2:1

Criss-cross the numbers and write them as subscripts.

THESE FORMULAS CANNOT CHANGE!!

B. Balance the equation using COEFFICIENTS. Each kind of atom should have the same

number on the left-hand side and the right-hand side of the equation (Law of Conservation of

Mass).

Remember to treat polyatomic ions as a unit (just like an element).

Treat H2O as H(OH) when balancing an equation.

i.e. H2SO4 + KOH → K2SO4 + H2O / H(OH)

Balance atoms that appear in their elemental form last.

Always balance oxygen and hydrogen last.

Check to see if the equation is balanced.

Problems: Write a balanced chemical equation for the following:

A. sodium metal + oxygen gas sodium oxide powder

B. lithium oxide lithium + oxygen

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AA. TYPES OF CHEMICAL EQUATIONS

Note: In the following general reactions the letters A, B C, and D will be used to represent

elements, compounds or polyatomic ions.

1. Synthesis Reaction (S): when simpler substances (usually elements) combine to form a

compound.

General reaction: A + B AB

Example #1: Mg + Cl2 MgCl2

Example #2: 2Na + S Na2S

Example #3: Al + N2 _______________________________________

2. Decomposition Reaction (D): when a compound decomposes into simpler substances

(usually elements).

General reaction: AB A + B (the opposite of synthesis)

Example #1: MgCl2 Mg + Cl2

Example #2: Na2S 2Na + S

Example #3: FeBr3 _______________________________________

3. Single Displacement Reaction (SD): when a compound reacts with an element.

If C is a metal, it replaces the metal A

General reaction: AB + C CB + A

Example #1: CuSO4 + Fe FeSO4 + Cu

Example #2: Ca(NO3)2 + 2Na 2NaNO3 + Ca

Example #3: MgCO3 + Rb _______________________________________

If C is a non-metal, it replaces the non-metal B

General reaction: AB + C AC + B

Example #1: 2NaBr + Cl2 2NaCl + Br2

Example #2: 2Na2S + O2 2Na2O + 2S

Example #3: FeCl3 + F2 _______________________________________

4. Double Displacement Reaction (DD): when a compound reacts with another compound.

Metals switch with metals OR non-metals switch with non-metals. Or positive ions

switch with positive ions OR negative ions switch with negative ions.

General reaction: AB + CD AD + CB

Example #1: 3CuS + 2FeCl3 Fe2S3 + 3CuCl2

Example #2: Ca(NO3)2 + 2NaOH 2NaNO3 + Ca(OH)2

Example #3: MgCO3 + RbCl _______________________________________

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5. Combustion (C): when a carbon compound reacts with oxygen, it will ALWAYS

produce carbon dioxide and water.

General reaction: carbon compound + O2 CO2 + H2O

Example #1: CH4 + 2O2 CO2 + 2H2O

Example #2: C2H5OH + 3O2 2CO2 + 3H2O

Example #3: C3H8 + O2 _______________________________________

Summary

Synthesis: element + element compound

Decomposition: compound element + element

Single displacement: element + compound element + compound

Double displacement: compound + compound compound + compound

Combustion: carbon compound + oxygen carbon dioxide + water

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BB. IB CHEMISTRY Y1 SUMMER ASSIGNMENT

SCIENTIFIC NOTATION

1. Express the following numbers in scientific notation. 42 000 000 000 ____________

2. Write out the following numbers in long form: 8.234 x 10-6

_____________

DIMENSIONAL ANALYSIS/ SI UNITS:

3. Identify the unit to the following mathematical operations:

a. 52.5 kg = 19 ?

2.8 kg

b. 56.5 g - 38.2 g = 0.22 ?

8.2 cm x 1.8 cm x 5.6 cm

c. 55.2 g + 26.8 g + 12.3 g = 94.3 ?

d. (20.5 mL)(1.00 g)(12.8 ºC) = 262 ?

mL

4. Convert the following measurement into the unit identified USING DIMENSIONAL

ANALYSIS.

a. Convert 112 cm to m.

b. Convert 21,510 mL to L.

c. Convert 2.18 kg to g.

d. How many liters are in 10.0 cm3?

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SIGNIFICANT FIGURES & CALCULATIONS

5. How many significant figures are there in 230. cm? _____________

6. How many significant figures are there in 0.05587 m ? _____________

7. Express the answer to each of the following calculations with the correct number of

significant figures.

a. 3.42 cm + 8.1 cm = _____________

b. 1.2 cm x 1.34 cm = _____________

8. Using dimensional analysis perform the following problems. Show all of your work. Pay

attention to significant figures.

a. If 1 kg = 2.205 pounds (lbs), how much does Sylvie weigh in pounds if she is 52.5 kg?

b. Would a car travelling at a constant speed of 65 km/h violate a 0.70 mi/ min speed limit?

Show your work. 1 mi = 1.6 093 km

46 | P a g e

CLASSIFICATION OF MATTER/ DEFINITIONS OF MATTER

9. Name one similarity and one difference between elements and compounds. Use the terms

sodium, chlorine and sodium chloride in your explanation. [4]

10. Classify A, C, E as either an element, compound, molecule, mixture of elements, mixture of

compounds, mixture of elements and compounds. [3]

A ___________________________________________________

C ___________________________________________________

E ___________________________________________________

A.

B. C. D.

E.

F. G. H.

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11. As above, draw an example at the atomic level of the following. [3]

A mixture of 1 element and

a compound

One molecule made of three

different types of atoms

A mixture of 2 different

compounds

THE MODERN VIEW OF ATOMIC STRUCTURE

12. Using the periodic table, fill in the following table. Assume the atom is neutral if no other

information is given.

Symbol Atomic

Number

Mass

Number

Charge Electrons Neutrons Protons

58Ni

2+

3919K

16 32 2-

17 37 18

3115P

3-

13 12

3 7 2

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13. Complete the table by describing the characteristics of each numbered element in the

periodic table below.

8

1

4 9

2 6

5 7

3

Number Group Period Class (metal,

non-metal,

metalloid)

Number of

valence

electrons

Outermost

energy level

1

2

3

4

5

6

7

8

9

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NAMING SIMPLE COMPOUNDS

14. Name the following compounds.

a. NaCl _______________________________________________

b. K2S _______________________________________________

c. CuBr2 _______________________________________________

d. FeF3 _______________________________________________

e. MgO _______________________________________________

f. NH3 _______________________________________________

g. SF6 _______________________________________________

h. Rb2SO4 _______________________________________________

i. AgClO3 _______________________________________________

j. Be3(PO4)2_______________________________________________

k. HNO3 _______________________________________________

l. H2SO4 _______________________________________________

m. HCl_______________________________________________

15. Write the formulas of the following compounds.

a. Sodium bromide __________________________________

b. Radium iodide __________________________________

c. Rubidium nitrite __________________________________

d. Magnesium oxide __________________________________

e. Francium sulfite __________________________________

f. Aluminum phosphide __________________________________

g. Carbon tetrachloride __________________________________

h. Dinitrogen pentoxide __________________________________

i. Terbium (IV) oxide __________________________________

j. Yttrium (III) nitride __________________________________

k. Titanium (IV) nitride __________________________________

l. Gold (I) oxide ___________________________________

m. Hydrogen sulfide __________________________________

n. Hydrogen phosphate __________________________________

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WRITING WORD & BALANCED CHEMICAL EQUATIONS

IDENTIFYING TYPES OF CHEMICAL EQUATIONS 16. Identify the type of equation, write the word equation and balanced chemical equation.

a. _____________ hydrogen gas and oxygen gas combine to produce water vapor

b. _____________ solid aluminum reacts with aqueous copper (II) chloride to produce solid

copper and aqueous aluminum chloride

c. _____________ lead (II) nitrate solution reacts with potassium iodide solution to

produce solid lead (II) iodide and aqueous potassium nitrate

d. _____________water breaks down into hydrogen gas and oxygen gas

e. _____________methane burns in oxygen to form carbon dioxide gas and water vapor

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17. FOR EACH OF THE FOLLOWING:

i) State the type of chemical reaction. (S, D, SD, DD, C)

ii) Complete the word equations.

iii) Write a balanced chemical equation.

a. ________copper (II) sulfate + barium chloride→

b. ________ethanol (C2H5OH) + oxygen →

c. ________calcium + aluminum sulfide →

d. ________aluminum bromide →

e. ________glucose + oxygen →

f. ________iron (II) chloride + potassium sulfate→

g. ________sodium + bromine →