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IB CHEMISTRY YEAR 1 SUMMER ASSIGNMENT

The IB Chemistry Year 1 Summer Assignment is a review of Chemistry from Grade 9 CPF (Chemistry & Physics Fundamentals) course at ASD. It will take between 3-5 hours for students who have taken CPF before. It will take longer for students who have no Chemistry background.

I. Read sections from Zumdahl Textbook II. Read class notes & complete sample problems

III. Complete IB Chemistry Y1 Summer Assignment (A test will be given in the first week of school- similar format to the assignment)

All material is available on pdf on www.asd.edu.qa Email any questions to jlamontagne@asd.edu.qa

Textbook Sections Class notes (with links)

Units & Measurement/ Significant Figures and Calculations

1. Scientific Notation 2. Dimensional Analysis 3. SI Units- The Metric System & Metric Conversion 4. Significant Figures 5. Significant Figures & Calculations

Molecules & Ions / The Modern View of Atomic Structure

1. Classification of Matter 2. Definitions to Classify Matter 3. Physical & Chemical Properties & Changes 4. The Modern View of Atomic Structure

An Introduction to the Periodic Table

1. The Modern Periodic Table 2. Parts of the Periodic Table 3. Using the Periodic Table 4. Chemical Symbols

 

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Naming Simple Compounds

1. Introduction to Naming Compounds 2. Naming Binary Ionic Compounds 3. Writing Formulas for Binary Ionic Compounds 4. Naming Ionic Compounds with Transition Metals 5. Writing Formulas for Ionic Compounds with Transition Metals 6. Naming Ionic Compounds with Polyatomic Ions 7. Writing Formulas for Ionic Compounds with Polyatomic Ions 8. Naming Binary Covalent Compounds 9. Writing Formulas for Binary Covalent Compounds 10. Naming Acids & Their Formulas 11. Chemical Nomenclature Flow Chart

Balancing Chemical Equations

1. Introduction to Chemical Equations 2. Rules for Writing Chemical Equations 3. Rules for Writing Balanced Chemical Equations 4. Types of Chemical Equations

 

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USEFULWEBSITES 

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 10bformat.

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 = ________mL2 or 1 m3 = ________________ cm3

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 = ______________________ cm3

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 atoms that are chemically combined.

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

Ionic compound

The atoms are different Ex: NaCl, NaHCO3 Compounds form large crystal lattice structures and the chemical formula

represents the simplest ratio. Usually made up of one metal and one non-metal atom

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

Molecular Compound

a pure substance made of two or more atoms that are chemically combined. The atoms can be the same O2, H2 The atoms can be different CO2, H2O Usually made up of all non-metal atoms

 

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5. Molecule: the smallest unit that makes up a 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, AZX. 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: 35Cl and 37Cl are isotopes. All elements have isotopes. Natural chlorine contains 75% 35Cl and 25% 37Cl. 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

65Cu2+

81Br1-

 

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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: Mg2+ Cl1- 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: NH4

1+CO32-

3. Reduce the oxidation numbers to their lowest possible ratio. Ex: NH41+CO3

2-

4. Put the polyatomic ions in brackets. (NH41+)(CO3

2-)

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. NAMING ACIDS & THEIR FORMULAS

1. When certain compounds are dissolved in water they produce a solution containing protons, H+. These substances are acids and will be discussed more extensively in its own topic, Acids & Bases.

2. Acid names are known by their common names rather than their formal name

(derived from the rules above). However, there are general rules to naming acids. 3. If the acid does not contain oxygen then the name takes the prefix hydro- and the

suffix –ic. i.e. HCl (aq) is hydrochloric acid, H2S (aq) is hydrosulfuric acid. 4. If the acid contains oxygen (in the anion), the acidic name is formed from the rrot

name of the anion with the suffix –ic or –ous depending on the name of the anion. If the anion name ends with –ate, then in the acid it is replaced by –ic. For

example H2SO4 contains the anion sulfate and it is called sulfuric acid. If the anion name ends with –ite, then in the acid it is replaced by –ous. For

example H2SO3 contains the anion sulfite and it is called sulfurous acid. Table: Name and formula of common acids that do not contain oxygen. Acid Formula Acid Name HF Hydrofluoric acid HCl Hydrochloric acid HBr Hydrobromic acid HI Hydroiodic acid HCN Hydrocyanic acid H2S Hydrosulfuric acid Table: Name and formulas of common oxygen-containing acids. Acid Formula Anion Acid Name HNO3 Nitrate Nitric acid HNO2 Nitrite Nitrous acid H2SO4 Sulfate Sulfuric acid H2SO3 Sulfite Sulfurous acid H2CO3 Carbonate Carbonic acid H3PO4 Phosphate Phosphoric acid HC2H3O2 Acetate Acetic acid HClO4 Perchlorate Perchloric acid HClO3 Chlorate Chloric acid HClO2 Chlorite Chlorous acid HClO Hypochlorite Hypochlorous acid

X. CHEMICAL NOMENCLATURE FLOW CHART

   

Y. 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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Z. 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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AA. 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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BB. 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

Decomposition: compound

Single displacement: element + compound

Double displacement: compound + compound

Combustion: carbon compound + oxygen

 

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CC. 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

 

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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.

Symbol Atomic Number

Mass Number

Charge Electrons Neutrons Protons

58Ni2+

3919K

16 -2

17 37 18

3115P

3-

13 12

3 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

3

7

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. Sulfurous acid __________________________________

n. Hydrosulfuric acid __________________________________

 

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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 vaopor

 

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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 →