Grade 10 PB Science

Chemistry UnitAtoms, the Periodic Table

andAn Introduction to Bonding and Chemical Reactions

1

Name: Day/Slot:

2

Contents

Anatomy of an Atom............................................................................................................................................. 3

Atomic Structure................................................................................................................................................... 4

Trends and Patterns in The Periodic Table............................................................................................................7

Using the Periodic Table......................................................................................................................................10

Quick Chemistry Review 1...................................................................................................................................11

Representing the Atom....................................................................................................................................... 12

Bohr & Lewis Dot Diagrams.................................................................................................................................14

Compare and Contrast: Bohr vs Lewis.................................................................................................................15

Ions and Electrons............................................................................................................................................... 18

Chemical Bonding................................................................................................................................................19

Ionic Bonds.......................................................................................................................................................... 21

Ionic Compounds: Names and Formulas.............................................................................................................26

Ionic bonding with transition metals...................................................................................................................27

Writing Ionic Formulas........................................................................................................................................ 30

Dice of Science.................................................................................................................................................... 31

Polyatomic Ions................................................................................................................................................... 32

Practice: Polyatomic ions.................................................................................................................................... 33

Assorted Ionic Bonding Questions.......................................................................................................................34

Covalent Compounds.......................................................................................................................................... 36

Review – Naming/Writing Chemical Compounds................................................................................................38

Counting Atoms................................................................................................................................................... 40

Investigation: The Law of Conservation Of Mass.................................................................................................43

Balancing Chemical Equations.............................................................................................................................44

Balancing equations practice...............................................................................................................................48

Balancing chemical equations more practice......................................................................................................48

Balancing chemical equations (word problems)..................................................................................................51

Chemical reaction types...................................................................................................................................... 52

Classifying Types of Chemical Reactions..............................................................................................................54

Identifying Types of Chemical Reactions.............................................................................................................57

Grade 10 IB Chemistry Unit Review.....................................................................................................................58

3

Anatomy of an Atom

4

5

Atomic Structure Date: Atom: a basic unit of matter that consists of a central nucleus (which is very dense) surrounded by a cloud of negatively charged electrons.

! One human hair is 1 million carbon atoms wide! One drop of water contains 2 sextillion (2x1021) atoms of oxygen, and 4 sextillion atoms of hydrogen

Element: a substance that cannot be divided into other chemical substances by a chemical process.

There are three particles in an atom, called subatomic particles:a. Protons:

- found in the nucleus

- has a positive charge

- not involved in bonding

- has a mass of slightly more than 1 amu or 1.6726 × 10-27 kg

- abbr. p+

b. Neutrons:- found in the nucleus

- has no charge (neutral)

- has mass of slightly more than 1 amu

- help hold the nucleus together

- some atoms of an element have different numbers of neutrons – these are called isotopes

- abbr. n0

Together the neutrons and protons form the nucleus.- very small dense region of an atom

- the size ranges from 1.6 fm to 15fm (1 fm = 1 x 10-15 m)

- overall positive charge

- makes up 1/100,000th % of the volume of the atom

- makes up 99.94% of the mass of the atom

c. Electrons:- found in 3D shells that orbit the nucleus (like Matryoshka doll)

- has a negative charge

- has almost no mass

- all bonding between atoms is based on electrons!

- unlike protons and neutrons, electrons cannot be broken down into smaller particles (i.e. they are not

made up of quarks)

- Abbr. e-

6

Atomic Mass: the sum of protons and neutrons (remember they each weight ~1 u)Example: Sodium (Na) has 11 p+ and 12 n0 atomic mass is 23

Atomic Number: the number of protons. - This number is unique to each element

- e.g. an atom with 8 protons is oxygen, and element with 36 protons is krypton

- the number of electrons is usually the same as the atomic number, but not always (i.e. ions!)

Atoms and electrons- The electrons that orbit around the nucleus do so in orbits, or shells.

- “shell” is a more appropriate term than “orbit”, because shell is 3-dimentional (like the layers of a jawbreaker)

- shells in an atom can hold up to eight electrons, except the first shell (i.e. the one closest to the nucleus)

- electrons cannot occupy an outer shell until the one closer to the nucleus is full

- atoms want to have their outer shell full of electrons – this is the bases for practically all of chemistry

- this outer shell is called the valence shell

- electrons in the valence shell are called valence electrons

Isotopes

- all atoms of the same element have the same number of protons (e.g. 6 protons always = carbon; 12 protons

always = calcium)

- the number of neutrons can vary (e.g. carbon can have 6 or 7 or 8 neutrons); this means the atomic mass of the

atom can be different for the same element (e.g. an atom of carbon with 6 neutrons would have an atomic mass

of 12 (6 + 6), a carbon atom with 7 neutrons would have an atomic mass of 13 (6 + 7))

- these different “versions” of elements are called isotopes

- isotopes are named a described based on their atomic mass (e.g. a carbon atom with 8 neutrons has an atomic

mass of 14, it is called carbon 14, and is abbreviated 14C)

- all elements have different isotopes, these different isotopes are the reason that atomic masses are not whole

numbers (e.g. carbon’s atomic mass is 12.01). The atomic mass is the average mass of all the atoms of that

sample, since some are heavier than others (because they have more neutrons)

- Try these:

Element Protons Neutrons Atomic Mass Isotope Name AbbreviationCarbon 6 7 13 Carbon 13 13C

Uranium 92 146 238 Uranium 238 238U

Lithium 3 3 6 Lithium 6 6Li

Magnesium 12 14 26 Magnesium 26 26Mg

Neon 10 12 22 Neon 22 22Ne

7

Use the information on the previous page to answer the following questions:1. Make a table below that lists the three parts of the atom. For each part, identify: its location; its abbreviation; its

relative mass (in atomic mass units); its charge.

2. Describe the shape of the space in which electrons exist.

3. Electrons account for what percentage of an atom’s mass?

4. If an atom were blown up to the size of gym A, how big do you think the nucleus would be? (Name an object that you think would be the same size as the nucleus)

5. An atom with an atomic mass of 2 could have how many protons? Explain

6. “Most of matter is empty space”. Explain this statement using information from the notes on the previous page.

7. Is there a difference between two hydrogen atoms joined by nuclear forces and a helium atom? Explain (hint: atomic number of hydrogen is 1, atomic number of helium is 2)

8. What is an isotope? What is the difference between carbon 12 and carbon 13?

A version of an element with a different number of neutrons12C = 6 neutrons13C = 7 neutrons

8

Trends and Patterns in The Periodic Table Date: - elements are listed in order of increasing atomic number

- elements with similar properties occur in columns which are called groups.

- the number above each group indicates how many electrons are in its valence shell

- rows are called periods and they are numbered beginning with 1 (this period only includes H and He)

- the period number corresponds with the number of electron shells

Group 1A (Alkali Metals) & Group 2A (Alkaline Earth Metals)- elements on the leftmost column (Group 1A) are called Alkali Metals (does not include hydrogen)

- elements in Group 2A are called Alkaline Earth Metals. They are taken from minerals (earth)

- the aqueous solutions of these groups are alkaline

- aqueous solution: a solution in which the solvent is dissolved in water

- alkaline: the solution is basic (i.e. not acidic)

- both these groups are very reactive and are found in nature only combined with other elements

compounds, never as free elements

The Transition Elements & Lanthanides and Actinides- the transition elements (aka transition metals) fill the middle of the periodic table (periods 4 – 7)

- most are found only in compounds (exceptions: gold, silver, platinum, copper, mercury)

- many have more than one possible charge and are named using roman numerals

- Lanthanides and Actinides fill the bottom two rows of the periodic table

- they are transition elements

- they are relatively rare and not are commercially important as the other transition elements

Metalloids

9

- most elements on the periodic table can be classified as metal or a non-metal

- metalloids are the few elements display characteristics of both metals and non-metals

- there are 8 metalloids:

Boron (B) Silicon (Si) Germanium (Ge) Arsenic (As) Antimony (Sb) Tellurium (Te) Polonium (Po) Astatine (At)

- these elements form a ‘bridge’ between the metals (left side of the PT) and the non-metals (right side)

Non-Metals in Groups 4A-6A

- these elements are the most abundant elements in the Earth’s crust and

atmosphere

- several of these elements are also important in our bodies (eg: C, N, and O)

Halogens (Group 7A)

- these elements are highly reactive and commonly react with alkali metals

and with other metals

- alkali metals have 1 electron in their valence shell and halogens have 7

electrons in their valance shell

- they both want eight in their outer shell – alkali metals lose one electron

and the halogens gain one

- the name comes from the Greek words for salt (hals) and genes (gens)…so… the halogens all form salts

- chemistry definition of salt is different from the colloquial definition (i.e. not just table salt)

- salts are formed when acids and bases are mixed and neutralize each other

Noble Gases (Group 8A)- elements on the far right of the periodic table

- least reactive of all elements: they have 8 electrons in their outermost

shell so they don’t need any more electrons to fill their shell

- very rare: were only discovered about 100 years ago

10

On one of your periodic tables, indicate the following (you will find them at the back of this booklet)

Use the sheet “Trends and Patterns in the Periodic Table” to help with this task

Please read the following instructions carefully: Groups (use arrows to indicate what a ‘group’ is) Periods (use arrows to indicate what a ‘period’ is)

For the items listed below, make a legend on the back of your table: Metals (outline all metals in one colour – do not shade) Metalloids (outline all metalloids in one colour – do not shade) Non-Metals (outline all non-metals in one colour – do not shade)

Alkali metals (outline and shade all elements in this group one colour) Alkaline Earth metals (outline and shade all elements in this group one

colour) Transition metals, Lanthanide series and Actinide series (outline and shade

all elements in this group one colour) Non-metals in groups 4A, 5A and 6A (outline and shade all elements in this

group one colour) Halogens (outline and shade all elements in this group one colour) Noble gases (outline and shade all elements in this group one colour)

11

Using the Periodic Table Date _______

A) Fill in the following table. You should know the names/symbols of these 22 elements! Symbol Name of Element Symbol Name of Element1. N Nitrogen 12. F Fluorine

2. S Sulfur 13. Fe Iron

3. O Oxygen 14. Cl Chlorine

4. C Carbon 15. K Potassium

5. H Hydrogen 16. P Phosphorous

6. He Helium 17. Al Aluminum

7. Ni Nickel 18. I Iodine

8. Ag Silver 19. Na Sodium

9. Cu Copper 20. Mg Magnesium

10. Au Gold 21. Ca Calcium

11. Hg Mercury 22. Pb Lead

B) Fill in the following table. Use the information shown in the legends above the periodic table. Element Name Symbol Atomic

NumberAtomic Mass

State of Matter at Room Temp1

Metal, Nonmetal,or Metalloid

Hydrogen H 1 1.008 Gas Nonmetal

Boron B 5 10.81 Solid Metalloid

Sodium Na 11 22.99 Solid Metal

Nitrogen N 7 14.007 Gas Nonmetal

Nickel Ni 28 58.693 Solid Metal

Magnesium Mg 12 24.305 Solid Metal

Copper Cu 29 63.546 Solid Metal

Aluminum Al 13 26.982 Solid Metal

Calcium Ca 20 40.078 Solid Metal

Iron Fe 26 55.845 Solid Metal C) On looseleaf, make a 4-column chart listing the 22 common elements from part A above, and their uses. This information can be found on the website: www.webelements.com . Click on the symbol of the element, then find the information on the page (for state at room temp look at “Standard State”)

Symbol Element State at Room Temp

Interesting Information

N Nitrogen Gas Nitrogen compounds are found in foods, fertilizers, poisons, and explosives.

Quick Chemistry Review 1

1 Use the website listed in Part C to help with this column

12

1. What does the atomic number of an atom represent?

2. Can the atomic number of an element change? Explain.

3. What is the valence shell? What are valence electrons?

4. Can the number of electrons in an atom change?

5. What does the atomic mass of an atom represent?

6. Can the number of neutrons in an atom change?

7. What is an isotope?

13

Representing the Atom Date:

There are two ways of representing the atom that we will use: the and (sometimes called Lewis Dot Diagrams)

Bohr ModelsThe Bohr model was described by Danish scientist . The Bohr model is also called the because the electrons orbit the nucleus in a manner similar to the planets orbiting the sun.

The Bohr model does NOT describe the , as the structure of the atom beyond the nucleus by cannot be defined.2 However it is a sufficient representation of the atom for our purposes.

Review: The nucleus is made of two types of particles: and . Unlike electrons, these particles have mass. Electrons exist in shells that ‘orbit’ around the nucleus.

To find the number of protons in an atom use the . Example: Ca =20

To find the number of neutrons in an atom the atomic mass, then the atomic number. Example: Cl atomic number = 17, atomic mass = 35.45 (round to 35), neutrons = 35 – 17 = 18

To find the number of electrons in an atom when it has a , use the atomic number. Example: Mg: 12

Symbol Name Protons Neutrons Electrons M/NM/M

Li

Phosphorous

18

35

H

The number of electrons that can exist in one shell is fixed, and follows the . This rule states that a maximum of can exist in each shell. The exception to this rule is the (the first shell), which can hold a maximum of .

Drawing Bohr diagrams:1. Find the number of p+,e-,n0 for the atom2. Draw a circle to represent the nucleus. Inside it, indicate the number of protons and neutrons3. Draw the electron orbits around the outside of the nucleus (this number corresponds to the period in which the

element is found). Example: hydrogen is in period so it has electron shells. Magnesium is in period so it has electron shells.

4. Fill in the electrons using the following rules:

2 Years after Bohr’s work it was found that electrons do not move in orbits/shells, but rather in regions of high and low probability

14

a. Think of the orbits like the face of a clock. We will only add electrons at 12, 3, 6 and 9 o’clock only.b. The inner-most shell has only two electrons, and they both go at 12 o’clock.c. Adding the electrons in the next shell(s) is done in the following sequence:

i. 12 o’clockii. 12 o’clock

iii. 3 o’clockiv. 6 o’clockv. 9 o’clock

vi. 3 o’clockvii. 6 o’clock

viii. 9 o’clock

Note the Bohr models can only be used to represent atoms of elements

Note the number of electrons in the outermost shell (valence) should be the same as the (roman numeral)

Examples

Hydrogen Sulphur

Lewis Dot StructuresLewis diagrams show only the electrons in the atom’s . The centre of a Lewis diagram shows the element’s , not the number of protons and neutrons

Hydrogen Sulphur

15

Bohr & Lewis Dot DiagramsDraw Bohr diagrams in the top half and Lewis dot diagrams in the bottom half

for the first 18 elementsAnswer the following questions:1. Complete the Compare and Contrast Frame on the back of this sheet.2. Do you notice any similarities going across or down the table?3. Which model do you think is easier to understand? 4. Where would a Bohr diagram be useful?5. What would a Lewis dot diagram be useful?

16

1 2

3 10987654

1514131211 181716

Compare and Contrast: Bohr vs Lewis Unit __________________________ Topic __________________

17

How are _______________________ and __________________________ alike?

How are _______________________ and __________________________ different?

Write a statement to compare and contrast the two terms, concepts, or events.

COMPARE

CONTRAST

Atoms With Charges Date:

Remember that the number of protons in an atom is (i.e. it does not change) but that the number of

electrons and neutrons . When atoms , they become charged particles

called . This happens because the number of positively charged particles (protons) and negatively

charged particles (electrons) are .

If the number of protons is greater than the number of electrons, the ion will have a ;

positively charged ions are called . If the number of electrons is greater than the number of

protons, the ion will have a ; negatively charged ions are called .

The number of determines whether the atom will gain or lose electrons. Atoms with

valence electrons (i.e. members of groups 1 – 13) will typically

, atoms with valence electrons (15 – 18) will typically

. The gaining or losing of electron occurs because atoms have a desire to

. Now complete the table below:

Particle Name Charge More p+/e- Gain/Lose e- Type of Element (metal/non-metal)

All atoms other than those of the noble gasses have a ‘desire’ to fill their valence shell. In other words, all elements on

the periodic table are ‘trying’ to be like the noble gases. Atoms do this through .

To determine the charge of an ion, first determine whether the atom will gain electrons or lose electrons (is it a metal

(lose) or a non-metal (gain)), then determine the number electron(s) that will be gained or lost (the group it’s in). The

number of electrons gained or lost is the charge of the ion.

Elements in group 1 will electron, they will have a charge of

Elements in group 2 will electrons, they will have a charge of

Elements in group 13 will electrons, they will have a charge of

Elements in group 14 will electrons, they will have a charge of

Elements in group 15 will electrons, they will have a charge of

Elements in group 16 will electrons, they will have a charge of

Elements in group 17 will electron, they will have a charge of

Elements in group 18 will electrons, they will have a charge of

18

Examples: Calcium Chlorine

Will it gain or lose e-? How do you know?

How many e- will be gained/lost? How know?

Writing Ion Symbols/Formulas

When writing the ‘symbol’ or ‘formula’ of an ion, we write the of the element, followed the

in superscript. We never write “1”s. If the charge is +1 we write “+”, if it is -1 we write “-”

Examples: Calcium = Ca2+ Potassium = K+ Flourine = Phosphorous =

Sulphur = Oxygen = Magnesium = Aluminium =

Naming Ions

Naming ions is easy. If it is a cation (i.e. a metal), we use the , followed by the word

. If it is an anion (i.e. a non-metal) we of the element’s name and add the

suffix .

Examples:

Lithium = lithium ion

Sulphur = sulphide.

Magnesium =

Hydrogen =

Complete the worksheet on the next page called “Ions and Electrons” using the information above.

19

Ions and Electrons Date:

Atomic Symbol

Protons Neutrons Neutral

# of e-Change

of e-New # of

e-Ion

Formula Ion Name

1. Cl 17 18 17 gains 1 e- 18 Cl- Chloride

2. K 19 20 19 loses 1 e- 18 K+ Potassium ion

3. S 16 16 16 gains 2 e- 18 S2- Sulfide

4. Mg 12 12 12 loses 2 e- 10 Mg2+Magnesium

ion

5. Al 13 14 13 loses 3 e- 10 Al+3 Aluminum ion

6. Br 35 45 35 gains 1 e- 36 Br- Bromide

7. O 8 8 8 gains 2 e- 10 O2- Oxide

8. P 15 16 15 gains 3 e- 18 P3- Phosphide

9. F 9 10 9 gains 1 e- 10 F- Fluoride

10. Se 34 45 34 gains 2 e- 36 Se2- Selenide

11. Na 11 12 11 loses 1 e- 10 Na+ Sodium ion

12. H 1 0 1lose/gain

1 e- 0/2 H+/-Hydrogen

ion/hydride

13. Te 52 76 52 gains 2 e- 54 Te2- Telluride

14. N 7 7 7 gains 3 e- 10 N+3 Nitride

15. Li 3 4 3 loses 1 e- 2 Li+ Lithium ion

16. Ca 20 20 20 loses 2 e- 18 Ca2+ Calcium ion

17. I 53 74 53 gains 1 e- 54 I- Iodide

18. Ba 56 81 56 loses 2 e- 54 Ba2+ Barium ion

19. F 9 10 9 gains 1 e- 10 F- Fluoride

20. Cs 55 78 55 loses 1 e- 54 Cs+ Caesium ion

Chemical Bonding Date: 20

Noble gases are located in the column on the Periodic Table (group 18/8A). These

elements are unique because they all have valance shells that are . All the other elements want

to have a full valence shell like the noble gases. There are three ways that elements can obtain a full valance

shell. They can:

1. valance electrons

2. valance electrons

3. valence electrons

It is a good thing for us that atoms want to fill their valence shell – without it the world as we know it would

not exist. Atoms try to fill their outer shells by combining with other atoms to form compounds. Compounds

are pure chemical substances that are made of two or more different

. Compounds often have properties

from the elements they are made of. For example, water is a

liquid at room temperature, but its constituents, oxygen and hydrogen, are both gases at room temperature.

There are two different ways that atoms can combine, or bond:

or

The type of bond depends on the type of elements that involved:

When a combines with a , they form an IONIC

BOND.

When a combines with a , they form a

COVALENT BOND.

*Two metals will bond to form a compound – if metals mix together, the substance

that is formed is called an .

21

On the chart below, indicate the type of bond that would be found in the following compounds (remember the

type of bond depends on the type of elements involved: metal or non-metal!). If the bond is ionic, circle the

element that is a metal.

Compound Ionic or Covalent Compound Ionic or Covalent

CaF2 Ionic FeS Ionic

NaCl Ionic NH3 Covalent

P2O5 Covalent CH4 Covalent

H2O Covalent MgO Ionic

LiH Ionic CO Covalent

CF4 Covalent H2S Covalent

Question In the compounds that had ionic bonds what did you notice about where the metal was in the

formula?

22

Ionic Bonds Date:

Ionic bonds occur when electrons are from one atom to another. When an atom

gains or loses electrons, it becomes a called an ion. Metals form ions called

(+), non-metals form ions called (–).

When two oppositely charged ions are near each other, they become attracted to each other like magnets -

this attraction is what causes ionic bonds.

Exercise: The illustrations below shows what happens during an ionic bond:

1. Identify the two elements on the right

2. How many valence electrons does the metal have?

3. How many valence electrons does the non-metal have?

4. How could these elements both get a full valence shell?

5. List two things that are different about each atom compared to the image above.

6. What do you think will happen to the positively charged Na ion and the negatively charged Cl anion?

23

Formulas for Ionic Compounds

Chemical formulas tell us the of each

element in a compound. For example, in the compound NaCl (sodium chloride), there is one atom of sodium

(Na) for each atom of chloride (Cl). In the compound MgF2 (magnesium fluoride), there is one atom of

magnesium (Mg) for every two atoms of fluoride (F).

Na2O

Exercise: Identify the number of atoms of each element in the ionic compounds listed below:

Compound Metal Number of atoms Non-Metal Number of atoms

Li3N (lithium nitride) Lithium 3 Nitrogen 1

K2O (potassium oxide) Potassium 2 Oxygen 1

BeF2 (beryllium fluoride) Beryllium 1 Fluorine 2

Al2S3 (aluminium sulphide) Aluminum 2 Sulphur 3

K2O (potassium oxide) Potassium 2 Oxide 1

AlBr3 (aluminium bromide) Aluminum 1 Bromide 3

CuCl2 (copper II chloride) Copper 1 Chlorine 2

Fe2S3 (iron III sulphide) Iron 2 Sulphur 3

Remember the goal of each atom in the compound is to

. Metals do this by electrons, non-metals do this by

electrons. Each atom in the compound MUST have a full valence shell. This often results in

different numbers of atoms of each element.

Example: Aluminum chloride

Draw a Lewis diagram:

24

How many chlorine atoms are needed for aluminium to give away all its valence electrons?

Example: Magnesium phosphide

Draw a Lewis diagram:

What ratio of magnesium to phosphorous atoms is needed to fill the valence shell of each atom?

Exercise: Determine the number of atoms of each element that would be needed to fill all valence shells.

Metal and Non-metal Atoms of Metal Atoms of Non-metal

Potassium and fluorine 1 1

Magnesium and chlorine

Lithium and nitrogen

Beryllium and oxygen

Aluminum and sulphur

25

Step-by-Step: If you are given the name of an ionic compound, to determine the formula of the compound do the following:

Steps Example – sodium oxide1. Write the ion for each element in the compound

(i.e. the symbol with the charge of the ion in the upper left)

Na+ O2-

2. Drop the charge from the symbol Na O

3. Write the charge of the opposite ion as the number of atoms of the other element (i.e. “criss-cross” the numbers) (never write “1”)

Na2O

Note: in step 3, if you end up with the same numbers as superscripts, reduce them. Example:

Mg2+ and O2- Mg2 and O2 Mg2O2 MgO

Exercise: Determine the correct formula for the following ionic compounds

Compound Name Formula Compound Name Formula

Sodium chloride NaCl Aluminium oxide Al2O3

Potassium sulphide K2S Beryllium bromide BeBr2

Calcium phosphide Ca3P2 Strontium iodide SrI2

Aluminium sulphide Al2S3 Sodium nitride Na3N

Lithium sulphide Li2S Barium phosphide Ba3P2

Naming Ionic Compounds

Naming ionic compounds is easy:

1. Write the name of the metal (don’t change anything)

2. Write the name of the anion (i.e. the non-metal with “ide” at the end)

List of common non-metals and their anions:

Non-metal Anion name Non-metal Anion nameBromine Bromide Nitrogen NitrideChlorine Chloride Oxygen OxideFluorine Fluoride Phosphorous PhosphideHydrogen Hydride Selenium SelenideIodine Iodide Sulphur Sulphide

26

Exercise: Determine the correct name for the following ionic compounds

Formula Compound Name Formula Compound Name

Li3P Lithium Phosphide NaF Sodium Fluoride

BeBr2 Beryllium Bromide Mg3N2 Magnesium Nitride

AlN Aluminum Nitride BaCl2 Barium Chloride

K2O Potassium Oxide Rb2Se Rubidium Selenide

SrI2 Strontium Iodide Cs3P Caesium Phosphide

Complete the worksheet “Ionic Compounds: Names and Formulas”. Ignore the following questions for now:

1) h, k, l, m, o, p, q, r, s, and 2) i, j, k, I, o, p, q, r

27

Ionic Compounds: Names and Formulas 1. Write the correct formula for the following ionic compounds

a) Magnesium oxide MgO k) copper (I) bromide CuBr

b) Sodium fluoride NaF l) tin (II) iodide SnI2

c) Aluminium nitride AlN m) iron (III) chloride FeCl3

d) Potassium sulphide K2S n) calcium phosphide Ca3P2

e) Lithium iodide LiI o) lead (II) oxide PbO

f) Calcium bromide CaBr2 p) lead (IV) fluoride PbF4

g) Beryllium oxide BeO q) tin (IV) bromide SnBr4

h) Nickel (I) chloride NiCl r) copper (II) sulphide CuS

i) Magnesium nitride Mg3N2 s) iron (II) oxide FeO

j) Aluminium sulphide Al2S3 t) calcium nitride Ca3N2

2. Write the correct name for each of the following ionic compounds

a) Li2O Lithium oxide k) PbS Lead II sulphide

b) AlCl3 Aluminum chloride l) SnO2 Tin IV oxide

c) MgS Magnesium sulphide m) Na2S Sodium sulphide

d) CaO Calcium oxide n) Mg3P2 Magnesium phosphide

e) KBr Potassium bromide o) NiO Nickel II oxide

f) BeF Beryllium fluoride p) CuI Copper I iodide

g) Na3N Sodium nitride q) PbCl4 Lead IV chloride

h) Al2O3 Aluminum oxide r) FeP Iron III phosphide

i) CuCl2 Copper II chloride s) CaF2 Calcium fluoride

j) FeBr3 Iron III bromide t) K3P Potassium phosphide

28

Ionic bonding with transition metals

ReviewHave you noticed that in all the ionic compounds we have worked with so far, the ?

For example: sodium chloride or NaCl is made of a (Na+) and a (Cl-). There is

, resulting in an overall charge of .

For example: aluminium oxide or Al2O3 is made of (Al3+) and (O2-). The total (2 × Al3+ = 6+), the total (3 × O2– = 6–). 6+ and 6- = 0.

Some Transition Metals Have Multiple ChargesSome metals are able to form more than one kind of ion (i.e. they can have ). Most of the transition metals – found in the – are able to form more than one type of ion.

For example: Copper (Cu) can form a cation with a charge of + (Cu+) or 2+ (Cu2+)Copper can react with chlorine to form 2 kinds of copper chloride:

white compound (Cu has 1+ charge, CuCl)yellow compound (Cu has 2+ charge, CuCl2)

Exceptions include , and .

Writing Formulas of Ionic Compounds with Transition MetalsWhen writing the formulas of compounds with transition metals the charge of the (i.e. the transition metal) is indicated using .

For example: nickel III bromide – the charge of

iron II phosphide – the charge of

titanium IV oxide – the charge of

For reference: roman numerals up to 10I – 1 III – 3 V – 5 VII – 7 IX – 9II – 2 IV – 4 VI – 6 VIII – 8 X – 10

29

For practice: determine the correct formula for the following ionic compounds:

platinum IV bromide

uranium VI selenide

iron III phosphide

lead IV nitride

manganese IV oxide

Naming Ionic Compounds with Transition MetalsIonic compounds with transition metals use roman numerals to indicate the charge of the cation. When given the formula for an ionic compound, the is to determine whether the cation (groups 3 – 12 plus (Tl), (Sn), (Pb), (Sb), (Bi) and (Po)).

If the cation IS a transition metal, its charge must be determined; this can be done by determining the of the anions.

Note that in all ionic compounds the positive charge and negative charges cancel each other out:

For example: Na3P – the charge of the cations is 3+ (3 × Na+), the charge of the anion is 3- (1 × P3-). 3+ and 3- = 0

For example: Al2S3 – the charge of the cations is 6+, the charge of the anions is 6-6+ and 6- = 0

This knowledge can be used to determine the charge transition metal cations:

For example: NiP – nickel is a transition metal, so its charge is unknown. The charge of all the anions is 3- (P3-). The charge of nickel must cancel the charge of the anion, so it must be 3+.

The name of this compound is

For example: ZnCl2 – the charge of the cation (zinc) is unknown because it is a transition metal. The charge of all the anions is 2- (2 × Cl-). The charge of the cation must cancel the charge of the anions, so it must be 2+.

The name of this compound is

For example: Co3As2 – cobalt is a transition metal, so its charge is unknown. The charge of all the anions is 6- (2 × As3-). The charge of cobalt must cancel the charge of arsenic, so it must be 6+. BUT WAIT! there are 3 atoms of cobalt in this compound, and the TOTAL charge of the cations is 6+, so EACH cobalt must be 2+ (6+ ÷ 3).

The name of this compound is .

30

Try these: 1. CuBr

What is the total charge of all anions?

What must the charge of all the cations be?

What is the charge of each cation?

What is the name of the compound?

2. Fe3P2

What is the total charge of all anions?

What must the charge of all the cations be?

What is the charge of each cation?

What is the name of the compound?

3. SnI2

What is the total charge of all anions?

What must the charge of all the cations be?

What is the charge of each cation?

What is the name of the compound?

4. PbO2

What is the total charge of all anions?

What must the charge of all the cations be?

What is the charge of each cation?

What is the name of the compound?

5. US2

What is the total charge of all anions?

What must the charge of all the cations be?

What is the charge of each cation?

What is the name of the compound?

31

Writing Ionic Formulas Date:

1. copper (II) hydride CuH2

2. sodium nitride Na3N

3. lithium oxide Li2O

4. cobalt (III) phosphide CoP

5. aluminum sulphide Al2S3

6. sodium nitride Na3N

7. iron (III) phosphide FeP

8. vanadium (V) oxide V2O5

9. sodium chloride NaCl

10. manganese (III) fluoride MnF3

11. beryllium nitride Be3N2

12. nickel (III) sulphide Ni2S3

13. potassium oxide K2O

14. silver bromide AgBr

15. zinc nitride Zn3N2

16. copper (II) arsenide Cu3As2

17. nickel (II) selenide NiSe

18. manganese (IV) oxide Mn2O

19. lead (IV) nitride Pb3N4

20. tin (II) hydride SnH2

21. lithium arsenide Li3As

22. chromium (VI) sulphide CrS3

23. calcium bromide CaBr2

24. lithium sulphide Li2S

25. copper (II) oxide CuO

26. platinum (IV) phosphide Pt3P4

27. aluminum telluride Al2Te3

28. silver nitride Ag3N

29. magnesium iodide MgI2

30. nickel (III) nitride NiN

31. vanadium (IV) phosphide V3P4

32. silver sulphide Ag2S

33. cobalt (III) sulphide Co2S3

34. iron (II) sulphide FeS

35. copper (II) nitride Cu3N2

36. nickel (II) oxide NiO

37. zinc nitride Zn3N2

38. manganese (VII) nitride Mn3N7

39. gallium sulphide Ga2S

40. ammonium selenide (NH4)2Se

32

Dice of Science

With a partner, you will create two dice of science – let the rolling begin!!!!!!!!!!!!

You will make a table to help keep track of your rolls. This is what your table should look like

Roll Balanced Formula Correct NameNa+ P3- Na3P Sodium phosphide

Yours will have 36 rows – one for each possible combination of the ions on your dice

33

Polyatomic Ions DATE:

A charged particle with more than one atom They are bonded covalently (i.e. electrons are shared, not exchanged) and are not neutrally charged –

they do not have an equal number of protons and electrons

What do they look like?

Ammonium Nitrite Nitrate

Ionic Compounds with Polyatomic Ions1. Formula: the same as ionic compounds (must be balanced – i.e. ‘criss-cross’)2. Naming: name of the metal followed by the name of the polyatomic ion (you do not need to change

the name of the polyatomic ion – i.e. no ‘ide’)

Examples1. sodium + nitrate

Na+ NO3- → NaNO3 sodium nitrate

2. sodium + sulfateNa+ SO4

2- → Na2SO4 sodium sulfate

3. barium + hydroxideBa2+ OH- → Ba(OH)2 barium hydroxide

34

Practice: Polyatomic ions

Write the formula. Write the name.

1. sodium chromate Na2CrO4 1. BeSO4 Beryllium sulfate

2. tin (IV) sulfate Sn(SO4)2 2. Pb(CN)4 Lead IV cyanide

3. aluminum hydroxide Al(OH)3 3. NH4NO3 Ammonium Nitrate

4. cesium cyanide CsCN 4. Zn(ClO3)2 Zinc II chlorate

5. beryllium chlorate Be(ClO3)2 5. K2Cr2O7 Potassium dichromate

6. iron (III) acetate Fe(C2H3O2)36. Mn2(Cr2O7)7

Manganese VII dichromate

7. calcium phosphate Ca3(PO4)2 7. HClO2 Hydrogen chlorite

8. ammonium sulfate (NH4)2SO4 8. FePO4 Iron III phosphate

9. barium nitrate Ba(NO3)2 9. Sr3(PO4)2 Strontium phosphate

10. calcium acetate Ca(C2H3O2)2 10. AuNO3 Gold I nitrate

11. sodium nitrate NaNO3 11. FeClO Iron I hypochlorite

12. copper (II) hydroxide Cu(OH)2 12. Ni(HSO4)3 Nickel III hydrogen sulfate

35

Assorted Ionic Bonding Questions Date:

Give the correct name of these ionic compounds

1. CuCl2 Copper II chloride 2. TiO3 Titanium VI oxide

3. BeO Beryllium oxide 4. CuF2 Copper II fluoride

5. Na2SO4 Sodium sulfate 6. Li2Cr2O7 Lithium dichromate

7. AlAs Aluminum arsenide 8. Pb(CN)2 Lead II cyanide

9. NH4NO3 Ammonium nitrate 10. Mg3(PO4)2 Magnesium phosphate

11. ZnO Zinc oxide 12. Tc(ClO3)2 Technetium II chlorate

13. Mg3N2 Magnesium nitride 14. (NH4)2O Ammonium oxide

15. Rb2O Rubidium oxide 16. CsCN Cesium cyanide

17. Fe(C2H3O2)3 Iron III acetate 18. BaF2 Barium fluoride

19. Ag3N Silver nitride 20. Au(NO3)2 Gold II nitrate

21. Ag2SO4 Silver sulfate 22. U(MnO4)6Uranium VI permanganate

23. Mn3N4 Manganese IV nitride 24. Np(OH)4Neptunium IV hydroxide

25. FeP Iron III phosphide 26. Pb(O2)2 Lead IV peroxide

36

Give the correct formula of these ionic compounds

1. Sodium chloride NaCl 2. Zinc oxide ZnO

3. Aluminium carbonate Al2(CO3)3 4. Tin (IV) selenide SnSe2

5. Calcium phosphate Ca3(PO4)26. Nickel (III) hydrogen

phosphate Ni2(HPO4)3

7. Potassium sulphide K2S 8. Gallium (III) peroxide Ga2(O2)3

9. Manganese (II) hydroxide Mn(OH)2 10. Iron (II) dichromate FeCr2O7

11. Gold (III) fluoride AuF3 12. Rubidium oxide Rb2O

13. Silver cyanide AgCN 14. Cobalt (III) nitride CoN

15. Beryllium fluoride BeF2 16. Lead (IV) phosphate Pb3(PO4)4

17. Francium phosphide Fr3P 18. Calcium chromate CaCrO4

19. Copper (II) iodide CuI2 20. Molybdenum (IV) nitride Mo3N4

21. Tungsten (V) telluride W2Te5 22. Mercury (I) chloride HgCl

23. Copper (I) acetate CuC2H3O2 24. Aluminium sulphate Al2(SO4)3

25. Radium arsenide Ra3As2 26. Uranium (VI) oxide UO3

37

Covalent Compounds Date:

Covalent compounds are formed between elements that are both non-metals. Covalent compounds are also called molecular compounds.

Unlike ionic compounds, where electrons are gained or lost, in covalent compounds the electrons are shared between atoms – this sharing of electrons IS the covalent bond.

Unlike ionic compounds where only one ratio is possible when elements bond (e.g. magnesium and fluorine can only bond with a ratio of 1 Mg to 2 F = MgF2), in covalent compounds multiple ratios between the same elements are possible:

o E.g. carbon and oxygen can form carbon monoxide (CO) and carbon dioxide (CO2)

Rules for covalent compounds1. Which element comes first?

The element closest to the upper left hand corner of the non-metals.

2. What do the prefixes mean?

The number of atoms of each element.

3. What are the prefixes?

Mono – 1Di – 2Tri – 3Tetra – 4

Penta – 5Hexa – 6Hepta – 7Octa – 8

Nona – 9Deca – 10

38

4. On which element does the ‘ide’ go?

The second element.

Food for thought… “Why don’t the names of ionic compounds include prefixes”

How to write formulas of covalent compounds1. Write the symbols of the two elements, followed by subscripts indicating the

number of atoms of each element. The number 1 is never written.Try these:a) Sulfur dioxide SO 2

b) Carbon tetrachloride CCl 4

c) Phosphorous pentafluoride PF 5

d) Dinitrogen pentoxide N 2O5

How to write names of covalent compounds

1. Write the names of the two elements; the first element is written in full, the second with “ide” on the end.

2. Add the correct prefix(es) to the names of each element. The prefix mono is only used with the second element, never with the first.

Try these:a) CO2 carbon oxide → Carbon dioxide

b) NO nitrogen oxide → Nitrogen monoxide

c) CCl4 Carbon tetrachloride

d) SF6 Sulfur hexafluoride

e) P2Br4 Diphosphorous tetrabromide

f) N2S Dinitrogen monosulfide

39

A. Name the following covalent compounds B. Write the formulas for the following covalent compounds

1) SeO2 Selenium dioxide

2) P2O5 Diphosphorus pentoxide

3) CI4 Carbon tetriodide

4) PBr3 Phosphorus tribromide

5) N2 Nitrogen gas

6) OCl2 Oxygen dichloride

7) NO3 Nitrogen trioxide

8) PCl5 Phosphorus pentachloride

9) SF6 Sulfur hexafluoride

10) C2Cl4 Dicarbon tetrachloride

1) Carbon monoxide CO

2) Nitrogen trichloride NCl 3

3) Tricarbon octahydride C 3H8

4) Selenium dibromide SeBr 2

5) Dinitrogen tetroxide N 2O4

6) Iodine monochloride ICl

7) Carbon tetrabromide CBr 4

8) Carbon disulphide CS 2

9) Boron trichloride BCl 3

10) Tetranitrogen pentafluoride N 4F5

40

Review – Naming/Writing Chemical Compounds Date: Formula Ionic or covalent Write the name of the compound.

1) NaBr Ionic Sodium bromide

2) CaO Ionic Calcium oxide

3) P2O5 Covalent Diphosphorus pentoxide

4) Ti(SO4)2 Ionic Titanium IV sulphate

5) FePO4 Ionic Iron III phosphate

6) K3N Ionic Potassium nitride

7) SO2 Covalent Sulphur dioxide

8) CuOH Ionic Copper I hydroxide

9) Zn(NO2)2 Ionic Zinc nitrite

10) V2S3 Ionic Vanadium III sulphide

11) NO2 Covalent Nitrogen dioxide

12) NaAt Ionic Sodium astetide

13) SiO2 Covalent Silicon dioxide

14) P2Br4 Covalent Diphosphorus tetrabromide

15) FeSO4 Ionic Iron II sulphate

16) SF6 Covalent Sulphur hexafluoride

17) Li2S Ionic Lithium sulphide

18) MgBr2 Ionic Magnesium bromide

19) N2S Covalent Dinitrogen monosulphide

20) Be(OH)2 Ionic Beryllium hydroxide

21) SO3 Covalent Sulphur trioxide

22) Cu2S Ionic Copper (I) sulphide

23) BF3 Covalent Boron trifluoride

24) CI2 Covalent Carbon diodide

25) Pb3(PO4)2 Ionic Lead II phosphate

Compound name Ionic or covalent Write the Formula of the compound.

41

26) silicon dioxide Covalent SiO2

27) nickel (III) sulfide Ionic Ni2S3

28) manganese (II) phosphate Ionic Mn3(PO4)2

29) silver acetate Ionic AgC2H3O2

30) diboron tetrabromide Covalent B2Br4

31) magnesium sulfate Ionic MgSO4

32) potassium carbonate Ionic K2CO3

33) ammonium oxide Ionic (NH4)2O

34) tin (IV) selenide Ionic SnSe2

35) carbon tetrachloride Covalent CCl4

36) carbon monosulfide Covalent CS

37) vanadium (II) phosphide Ionic V3P2

38) oxygen difluoride Covalent OF2

39) gold (I) phosphate Ionic Au3PO4

40) triboron tetrahydride Covalent B3H4

41) aluminum carbonate Ionic Al2(CO3)3

42) dinitrogen heptoxide Covalent N2O7

43) dinitrogen trioxide Covalent N2O3

44) cadmium (I) chloride Ionic CdCl

45) aluminum oxide Ionic Al2O3

46) disulfur trichloride Covalent S2Cl3

47) cobalt (II) acetate Ionic Co(C2H3O2)2

48) ammonium cyanide Ionic NH4CN

49) pentaphosphorus hexafluoride Covalent P5F6

50) hydrogen trisulfide Covalent HS3

Counting Atoms Date:

42

When dealing with chemical formulas and equations, it is important that you are able to COUNT the atoms involved in a certain compound or molecule.

Before we begin, let’s review some key terms:

Symbol: a letter or letters used to represent an element (e.g. H or Zn)

Atom: the smallest unit of matter that retains the unique properties of that element

Formula: a group of symbols and subscripts that are an abbreviation for a certain compound. (e.g. CO2 is the formula for carbon dioxide)

Subscript: a small number, located to the bottom right of a symbol, which indicates how many ATOMS there are of a certain element. (e.g. in CO2, the subscript is the 2. It means that there are 2 atoms of oxygen.)

Molecule: two or more atoms chemically bonded together. (e.g. H2O is the make-up of one molecule of water)

Coefficient: a large number, located to the left of a compound formula, which indicated how many MOLECULES there are of a certain compound. (e.g. in 6CO2, the coefficient is 6. It means there are 6 molecules of carbon dioxide.)

Use the terms above to label the following example:

How to count atoms in a compound43

3H2OThe number 3 is an example of a COEFFICIENT . It means that there are 3 molecules of H2O.

The number 2 is an example of a SUBSCRIPT . It means that there are 2 atoms of HYDROGEN in one molecule of WATER.

H2O is the FORMULA for the compound called WATER.

The O is the SYMBOL for the ELEMENT called oxygen.

The H is the SYMBOL for the ELEMENT called hydrogen.

1. SUBSCRIPTS If there is a subscript, it refers to the symbol on its left. The number means how many atoms of that element are in that compound.

For example, in C6H12O6 (glucose), there are 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. Every glucose molecule has exactly this many atoms – a total of 24 atoms in every glucose molecule.

PRACTICE: a) In one molecule of H2SO4 (sulfuric acid), there are 2 atoms of HYDROGEN, 1 atom of SULPHUR, and 4 atoms of OXYGEN.

b) In one molecule of AlCl3 (aluminum chloride), there is 1 atom of ALUMINUM and 3 atoms of CHLORINE.

44

2. BRACKETS If there is a polyatomic ion in brackets, the subscript on the outside of the brackets is MULTIPLIED by all of the atoms inside the parenthesis.

For example, in Ca(OH)2 (called calcium hydroxide), there is 1 atom of calcium. The subscript “2” gets multiplied by everything inside the brackets (hydroxide ion), so there are 2 atoms of oxygen, and 2 atoms of hydrogen.

Another example: Al2(SO4)3 (aluminium sulfate) there are 2 aluminium atoms. The 3 at the end does NOT affect the aluminum – it ONLY belongs to the sulphate. Because the 3 is referring to the whole polyatomic ion of sulphate (i.e. there are 3 sulphate ions), you will MULTIPLY the 3 by everything inside the brackets. SO: There are 3 sulphur atoms and 12 oxygen atoms.

45

PRACTICE:

a) In one molecule of Ba(OH)2 (barium hydroxide), there is 1 atom of BARIUM, 2 atoms of OXYGEN, and 2 atoms of HYDROGEN.

b) In one molecule of Ca(NO3)2, (calcium nitrate), there is 1 atom of CALCIUM, 2 atoms of NITROGEN, and 6 atoms of OXYGEN.

c) In (NH4)2SO4 (ammonium sulfate), there are 2 atoms of NITROGEN, 8 atoms of HYDROGEN, 1 atom of SULPHUR, and 4 atoms of OXYGEN.

3. COEFFICIENTS If there is a coefficient on the left of the formula, that number is MULTIPLIED by EVERYTHING on the right. BE SURE to figure out all the atoms with their subscripts and brackets FIRST, and then multiply the coefficient last.

46

For example, 6H2O means you have 6 molecules of water. Every water molecule has 2 hydrogen atoms and 1 oxygen atom. When you multiply everything by 6, you will get a TOTAL of 12 hydrogen atoms and 6 oxygen atoms.

Another example: In Pb(NO3)2 there is 1 atom of lead, 2 atoms of nitrogen, and 6 of oxygen. If we were to be given 2 Pb(NO3)2, however (note the coefficient of 2!) we would have to then multiply everything by 2. This gives us 2 atoms of lead, 4 atoms of nitrogen, and 12 atoms of oxygen in total.

PRACTICE:

a) In 3HNO3 there are 3 atoms of HYDROGEN, 3 atoms of NITROGEN, and 9 atoms of OXYGEN in total.

b) In 2Al2(SO4)3, there are 4 atoms of ALUMINUM, 6 atoms of SULPHUR, and 24 atoms of OXYGEN.

47

For each formula, write the number of atoms of each element. Compound TOTAL Number of atoms per element Compound TOTAL Number of atoms per element

a) AlCl3 1 Al, 3 Cl i) 2Pb(NO3)2 2 Pb, 4 N, 12 Ob) Pb(OH)2 1 Pb, 2 O, 2 H j) 2(NH4)2S 4 N, 16 H, 2 Sc) Al2(SO4)3 2 Al, 3 S, 12 O k) 5Al(OH)3 5 Al, 15 O, 15 Hd) 3CS2 3 C, 6 S l) 2C12H22O11 24 C, 44 H, 22 Oe) 2HNO3 2 H, 2 N, 6 O m) 10NH4Cl 10 N, 40 H, 10 Clf) 2Na2SO4 4 Na, 2 S, 8 O n) Ca(NO3)2 1 Ca, 2 N, 6 Og) 4Ca(OH)2 4 Ca, 8 O, 8 H o) 3H2SO4 6 H, 3 S, 12 Oh) 3(NH4)2SO4 6 N, 24 H, 3 S, 12 O p) 3Al(NO3)3 3 Al, 9 N, 27 O

48

Balancing Chemical Equations Date:

Often, chemical equations are described in words; for example, the burning of methane.

We have also seen that chemical reactions can be written as skeleton equations. For example:

However, there is a problem: the Law of Conservation of Mass says that in a chemical reaction, matter can be neither created nor destroyed. In other words, the atoms on the left (reactants) must equal the atoms on the right (products).

If we make a table of the number of atoms of each element, we see they are not equal:Element # Atoms Reactants # Atoms Products

C 1 1

H 4 2

O 2 3

It seems we have ‘created’ an oxygen atom, and ‘destroyed’ two hydrogen atoms.

We cannot change the types of elements or the formulas of compounds.

How can we correct this imbalance? The answer is to change the number of molecules, rather than their formulas.

If we add an oxygen molecule to the reactants, and add a water molecule to the products, this balances the equation:

Element # Atoms Reactants # Atoms Products

C 1 1

H 4 4

O 4 4

49

Steps to balance a chemical equation:Example: Iron reacts with oxygen to form iron oxide (Fe2O3)1. Write word equation for the reaction

Iron + Oxygen → Iron oxide

2. Write the equation by replacing names with formulas

Fe + O2 Fe2O3

3. Count the number of atoms of each type of element in the reactants and products

Element # Atoms Reactants # Atoms Products

Fe 1 2O 2 3

Hint when getting started - choose an element that appears only once on each side

4. Determine the correct coefficients to balance the equation (trial & error)

4 Fe + 3 O2 2 Fe2O3

A simpler way of keeping track of atoms is by writing the atoms in the reaction below the arrow, as done below:

50

1- Mg -1 2- H -22- N -26- O -6

Practice:1. Try balancing the following equations:

a) Na + Cl2 NaCl

2Na + Cl2 2NaCl

b) CO + O2 CO2

2CO + O2 2CO2

c) K + O2 K2O

4K + O2 2K2O

d) Al + Br2 AlBr3

2Al + 3Br2 2AlBr3

e) H2 + O2 H2O

2H2 + O2 2H2O

f) N2H4 + O2 H2O + N2

N2H4 + O2 2H2O + N2

g) H2 + Cl2 HCl

H2 + Cl2 2HCl

h) CH4 + O2 CO2 + H2O

CH4 + 2O2 CO2 + 2H2O

i) N2 + H2 NH3

N2 + 3H2 2NH3

2. For each of the following, write the correct skeleton equation, and then balance it to form a chemical equation.

a. Copper (II) oxide + hydrogen gas copper + water

CuO + H2 Cu + H2O (skeleton equation IS balanced equation)

b. Lead (II) nitrate + potassium iodide lead (II) iodide + potassium nitrate51

Skeleton Eq: Pb(NO3)2 + KI PbI2 + KNO3

Balanced Eq: Pb(NO3)2 + 2KI PbI2 + 2KNO3c. Calcium + water calcium hydroxide + hydrogen gas

Skeleton Eq: Ca + H2O Ca(OH)2 + H2

Balanced Eq: Ca + 2H2O Ca(OH)2 + H2

d. Lead (II) sulphide + oxygen lead + sulphur dioxide

PbS + O2 Pb + SO2

(skeleton equation IS balanced equation)

e. Hydrogen sulphide hydrogen + sulphur

H2S H2 + S(skeleton equation IS balanced equation)

3. You are an engineer trying to determine how much air must be supplied to burn gasoline in an engine. Assume gasoline is heptane (C7H16); the word equation is:

Heptane + oxygen carbon dioxide + water vapour

a. Write the skeleton equation for the reaction

C7H16 + O2 CO2 + H2O52

b. Balance the equation

C7H16 + 11O2 7CO2 + 8H2O

c. How many molecules of oxygen are required for every molecule of heptane that burns?

11

53

Balancing equations practice

54

3 4522 3

2222

6 6 6

74 92 32

34 1232 6

34284 45

3

6

2

8

3

316

23

22

24

3326

3 2

362

Balancing chemical equations more practice1. Cu + H2O CuO + H2

2. 2CO + O2 2CO2

3. 2KNO3 2KNO2 + O2

4. 2Fe + 3H2SO4 Fe2(SO4)3 + 3H2

5. 3O2 + CS2 CO2 + 2SO2

6. Cu + Cl2 CuCl2

7. 3Mg + N2 Mg3N2

8. C + O2 CO2

9. P4O10 + 6H2O 4H3PO4

10. 2K + 2H2O 2KOH + H2

11. 2NaOH + H2SO4 Na2SO4 + 2H2O

12. 2Al + 3H2SO4 Al2(SO4)3 + 3H2

13. NH4NO2 N2 + 2H2O

14. 2NH3 + 3CuO 3Cu + N2 + 3H2O

15. 2C2H6 + 7O2 6H2O + 4CO2

16. 2H2 + O2 2H2O

17. H3PO4 + 3KOH K3PO4 + 3H2O

55

18. 6K + B2O3 3K2O + 2B

19. HCl + NaOH NaCl + H2O

20. 10Na + 2NaNO3 6Na2O + N2

21. 4C + S8 4CS2

22. 2Na + O2 Na2O2

23. 2N2 + 5O2 2N2O5

24. 2H3PO4 + 3Mg(OH)2 Mg3(PO4)2 + 6H2O

25. 2NaOH + H2CO3 Na2CO3 + 2H2O

26. KOH + HBr KBr + H2O

27. 4Na + O2 2Na2O

28. 2Al(OH)3 + 3H2CO3 Al2(CO3)3 + 6H2O

29. 6Cs + N2 2Cs3N

30. Mg + Cl2 MgCl2

31. 10Rb + 2RbNO3 6Rb2O + N2

32. N2 + 3H2 2NH3

33. Al(OH)3 + 3HBr AlBr3 + 3H2O

56

Balancing chemical equations (word problems) Date:

1. Aluminum reacts with oxygen to form aluminum oxide

Al + O2 Al2O3

4Al + 3O2 2Al2O3

2. Ethane (C2H6) reacts with oxygen gas to form carbon dioxide and water vapour.

C2H6 + O2 CO2 + H2O

2C2H6 + 7O2 4CO2 + 6H2O

3. Potassium oxide reacts with water to form potassium hydroxide

K2O + H2O KOH

K2O + H2O 2KOH

4. Iron III oxide reacts with lithium sulphate to form iron III sulphate and lithium oxide

Fe2O3 + Li2SO4 Fe2(SO4)3 + Li2O

Fe2O3 + 3Li2SO4 Fe2(SO4)3 + 3Li2O

5. Ammonium carbonate breaks down (decomposes) into ammonia (NH3), carbon dioxide and water

57

(NH4)2CO3 NH3 + CO2 + H2O

(NH4)2CO3 2NH3 + CO2 + H2O

6. Calcium chloride reacts with silver nitrate to form silver chloride and calcium nitrate

CaCl2 + AgNO3 AgCl + Ca(NO3)2

CaCl2 + 2AgNO3 2AgCl + Ca(NO3)2

58

Chemical reaction types Now that you know how to balance chemical equations, you can learn about the different types of reactions.

Each reaction type has a standard equation. A standard equation usually uses the letters A, B, C and D to represent what

happens in a chemical reaction. In the example below, chemical A combines with chemical B to form a new substance:

AB.

e.g.

Although other types of reactions can occur, we will talk about five types of reactions. These five types of

reactions cover most of the chemical reactions that occur in nature.

For each type of reaction, you should be able to: (1) describe in words what happens; (2) know the standard

equation; and (3) be able to identify the reaction type from a chemical equation.

Synthesis

Synthesis reactions have two or more reactants which combine to form one product.

The standard equation of a synthesis reaction is:

An example of a synthesis reaction is:

Decomposition

Decomposition reactions have one reactant which breaks into two or more products.

The standard equation of a decomposition reaction is:

An example of a decomposition reaction is:

Single Displacement

59

The Five Types of Chemical Reactions

SynthesisDecomposition

Single DisplacementDouble Displacement

Combustion

Note: you have probably noticed that both the example equations are balanced. This will not always be the case. When you have mastered this topic, you will be able to identify and balance chemical equations.

In single displacement reactions, one part of one reactant separates and joins another reactant. Single displacement

reactions have two or more reactants and two or more products.

The standard equation of a single displacement reaction is:

A + BC C + BA

Note: what has happened here? Answer: C has been displaced by A. Now A is bonded to B and C is alone.

An example of a single displacement reaction is:

Double Displacement

In double displacement reactions, the positive ion of one reactant displaces the positive ion of the other reactant.

The two ions that have been left, now combine to form a new compound. Double displacement reactions have two

or more reactants and two or more products.

The standard equation of a double displacement reaction is:

AB + CD AD + CB

Note: what has happened here? Answer: D has displaced B, and B has displaced D. After the reaction (i.e. the

arrow) A is now bonded to D, and C is bonded to B. There are two displacements (B displaced D and D displaced

B) i.e. double displacement!

An example of a double displacement reaction is:

60

Combustion (Everyone’s favourite!)

Combustion reactions involve any reactant and oxygen. The products always include CO2 and water. Energy is

always given off in the form of heat.

The standard equation of a combustion reaction is

Note: what has happened here? Answer: Some substance (called ??) has combined with oxygen to form carbon

dioxide and water. Carbon dioxide and water will always be the products of a combustion reaction. Carbon

dioxide and water will always be the products of a combustion reaction. Carbon dioxide and water will always be

the products of a combustion reaction. Get it?

An example of a combustion reaction is:

61

Classifying Types of Chemical Reactions

Objectives:By the end of this worksheet you should be able to:

1. Know the names of the five types of reactions;2. Know the standard chemical equation for each of the five types of reactions; and3. Given a chemical equation, identity which type it is.

Types of chemical reactions1. List the five types of chemical reactions:

a. Synthesisb. Decompositionc. Single Displacementd. Double Displacemente. Combustion

2. Choose three of the reaction types and identify how many products and how many reactants are necessary for this type of reaction to take place.

Type of Reaction Number of Products Number of ReactantsDecomposition 2 1Synthesis 1 2Single Displacement 2 2Double Displacement 2 2Combustion 2 2

3. Using the letters A, B, C and D write the standard chemical equation for the two reaction types you did not use in question 2, plus one more.

Type of Reaction Standard EquationSynthesis A + B → ABDecomposition AB → A + BSingle Displacement A + BC → AB + CDouble Displacement AB + CD → AD + BCCombustion ?? + O2 → CO2 + H2O + Energy

4. Synthesis and decompositiona. The prefixes “synth” and “decomp” make up parts of many words. For each of these prefixes, identify

four words other than synthesis or decomposition that contain these prefixes.Two words that start with synth Two words that starts with decomp

b. Based on your words above, what do you think these prefixes mean?Synthesis Decomposition

62

c. Explain how these prefixes relate to the type of chemical reaction they describe:Synth Decomp

5. Fill in the following compare/contrast chart (you need a minimum of 3 points in each box!)How are single displacement and double displacement reactions different?

How are single displacement and double displacement reactions alike?

Write a two sentence statement to compare and contrast single and double displacement reactions

6. If you are given a chemical equation you must be able to distinguish a double displacement reaction from a single displacement reaction. Give two ways you can tell the difference based only on a chemical equation.

a.

b.

7. The only type of chemical reaction that requires oxygen is .a. List five combustion reactions from everyday life:

i. _____________________________________________________ii. _____________________________________________________

iii. _____________________________________________________iv. _____________________________________________________v. _____________________________________________________

8. Describe how you will distinguish each type of chemical equation from the othersReaction Type How you will distinguish it from other types

63

9. Match the type of equation on the left with the equation on the rightReaction Type EquationDouble displacement

Combustion

Single displacement

Synthesis

Decomposition

10. Go back to the list of “objectives” at the beginning of this sheet. Check-off the objectives you feel you understand. If you don’t feel confident with these, read the sheet over once more, then see your teacher.

11. On the sheet “Balancing Equations Race” identify each reaction by type.

64

Identifying Types of Chemical Reactions Date:

Choose the correct symbol for the type of reaction. Place that answer in the blank at the beginning of each equation and then balance each equation correctly.

S = synthesis SD = single displacementD = decomposition DR = double displacementC = combustion

1. SD 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

2. S 2CO(g) + O2(g) 2CO2(g)

3. DR FeS(s) + 2HCl(aq) FeCl2(aq) + H2S(g)

4. D 2NaNO3(s) 2NaNO2(s) + O2(g)

5. C CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

6. SD Fe(s) + 2CuNO3(aq) 2Cu(s) + Fe(NO3)2(aq)

7. SD 2KI(aq) + Cl2(g) 2KCl(aq) +I2(aq)

8. S 2Al(s) + 3S(s) Al2S3(s)

9. D 2KClO3(s) 2KCl(s) + 3O2(g)

10. C 2C4H10(g) + 13O2(g) 8CO2(g) + 10H2O(l)

11. SD WO3 + 3H2 W + 3H2O

12. DR PdCl2 + HNO3 Pd(NO3)2 + HCl

13. S Si + Cl2 SiCl4

14. DR RbBr + AgCl AgBr + RbCl

15. S CO + O2 CO2

16. D C6H12O6 C2H5OH + CO2

17. S NO + O2 NO2

18. C C3H8 + O2 CO2 + H2O

19. DR CaF2 + H2SO4 CaSO4 + HF

20. DR BaF2 + LiBr BaBr2 + 2LiF

21. S Fe + O2 Fe2O3

22. D MgO Mg + O2

65

Grade 10 IB Chemistry Unit Review Date: To complete this sheet you will need your periodic table and list of polyatomic ions

Part 1 – The Structure of the Atom

1. List the three sub-atomic particles. After the name of each particle, indicate its charge in brackets.

Electron ( - ) Proton ( + ) Neutron ( 0 )

2. In the space below, draw a picture of an atom of lithium (draw the neutrons and protons). Be sure to include the correct numbers of subatomic particles and have them in the correct location.

3. Of the three subatomic particles, which has the most mass? Protons and neutrons

4. Of the three subatomic particles, which has the least mass? Electrons

5. What is the name of the outer shell of an atom? Valence shell

6. What are electrons in this shell called? Valence electrons

7. Why are these electrons more important than the others? They are the only particles involved in chemical bonding

Part 2 – Using the Periodic Table

1. Identify each of the letters in the symbol below (i.e. what information goes in each letter’s location?):

2. What are “groups” on the periodic table? Columns

3. Give three examples of groups (i.e. their names) Alkali metals, alkaline earth metals, halogens

4. What do all elements in a group have in common? Same number of valence electrons

5. What are “periods” on the periodic table? Rows

6. What do all elements in a period have in common? Same number of electron shells

66

A Element symbol

W Isotope number

X Atomic number

Y Atomic mass

w

7. Identify the element that is in:

a. Group 1 period 4 Potassium (K)

b. Group 13 period 5 Indium (In)

c. Group 2 period 7 Radium (Ra)

8. Give examples of two elements that are in each of these categories. You MAY NOT use the same element twice:a. Metals Potassium Cobalt

b. Non-metals: Phosphorus Oxygen

c. Metalloids: Boron Silicon

d. Alkali metals: Lithium Sodium

e. Alkaline Earth metals: Beryllium Magnesium

f. Transition metals: Scandium Iron

g. Halogens: Fluorine Chlorine

h. Noble gases: Helium Neon

9. Complete the table below;Element

name Symbol Atomic number

Atomic mass

Metal, nonmetal,or metalloid Protons Neutrons Electrons Valence

electronsRubidium Rb 37 85.47 Metal 37 48 37 1Selenium Se 34 78.96 Non-metal 34 45 34 6Bromine Br 35 79.90 Non-metal 35 45 35 7Calcium Ca 20 40.08 Metal 20 20 20 2

10. What is an isotope?A different “version” of an element where the number of neutrons are different. Example carbon 12, carbon 13 and carbon 14 are different isotopes of carbon.

Part 3 – Atoms, Ions and Bonding1. List two differences and two similarities between Bohr and Lewis models of the atom:

Bohr Lewis

Differences

Show all electrons

Shows protons and neutrons

Does not show element symbol

Show only valence electrons

Does not show protons and neutrons

Shows element symbol

SimilaritiesShow valence electrons

Model of the atom

Show valence electrons

Model of the atom

67

Other answers are possible

2. Draw a Bohr model and Lewis diagram for each of the following three elementsNitrogen Aluminium Argon

Bohr model

Lewis diagram

3. What is an ion? An atom with an unequal number of protons and electrons (a charged particle)

4. What is the difference between a cation and an anion? Cation has a positive charge, anion has a negative charge

5. Complete the table below for anions and cationsParticle Name Charge More p+/e- Gain/Lose e- Type of Element (metal/non-metal)

Cation + Proton Lose Metal

Anion - Electrons Gain Non-metal

6. Complete the table below by indicating the ion symbol and the name of the ion for these six elementsElements Ion symbol Ion name Elements Ion symbol Ion name

Sodium Na+ Sodium ion Phosphorous P3- Phosphide

Selenium Se2- Selenide Aluminium Al3+ Aluminium ion

Calcium Ca2+ Calcium ion Bromine Br- Bromide

7. List three differences between ionic and covalent compounds/bonds (differences in naming do not count)

I: electrons are exchanged/gained and lost; C: electrons are shared

I: involves metals and non-metals; C: 2 non-metals

I: ions involved; C: ions not involved

8. Complete the table below by identifying whether the compound is covalent or ionic68

7p+

7n013p+

14n018p+

22n0

N Al Ar

Compound Ionic or Covalent Compound Ionic or Covalent

SiO4 C WS3 I

SrSe I Li3N I

CaI2 I C4H10 C

9. For transition elements:a. Where are they found on the periodic table? In the middle of the table

b. What kind of elements are they? Metals

c. What does the Roman numeral next to a transition element indicate? The charge of the metal

10. Complete the table below by indicating the Arabic numeral next to the roman numeral (note the numbers are NOT in order):VI 6 I 1 VII 7 III 3 VIII 8

IX 9 V 5 II 2 X 10 IV 4

11. Why do we use prefixes with covalent compounds and not with ionic compounds? Ions of two elements can

combine in only one ratio so indicating the number of atoms of each element is unnecessary. Two non-metal

elements can combine in multiple ratios, therefore we need to indicate the number of atoms of each element

12. Complete the table below by indicating the correct prefix for each of these numbers:1 Mono 2 Di 3 Tri 4 Tetra 5 Penta

6 Hexa 7 Hepta/septa 8 Oct 9 Nona 10 Deca

13. For polyatomic ions:a. What does the name “polyatomic ion” mean? Ions with more than one atom

b. Why do they have a charge? Because they have an unequal number of protons and electrons

14. Complete the table below by first indicating whether the compound in covalent or ionic, and then giving the correct formula or name.

Name/Formula Covalent or Ionic Name/Formulaa) Zinc II oxide I ZnO

b) Silicon dioxide C SiO2

c) Na2O I Sodium oxide

d) Magnesium nitride I Mg3N2

e) B2Br4 C Diboron tetrabromide

69

f) SnSe2 I Tin IV selenide

g) Manganese IV oxide I MnO2

h) CCl4 C Carbon tetrachloride

i) V2S3 I Vanadium III sulphide

j) Ni(H2PO4)2 I Nickel II dihydrogen phosphate

k) Gold I nitrite I AuNO2

l) B3H4 C Triboron tetrahydride

m) Nitrogen nonaoxide C NO9

n) Si2F3 C Disilicon trifluoride

o) Ammonium oxide I (NH4)2O

p) OF2 C Oxygen difluoride

q) W(SO4)3 I Tungsten VI sulphate

r) Carbon diodide C CI2

s) Ca3(PO4)2 I Calcium phosphate

Part 4 – Chemical Reactions

1. Complete the table below by indicating the correct number of atoms of each element:Atoms of Each Element

5NH4Cl N – 5; H – 20; Cl – 5

W2(SO4)3 W – 2; S – 3; O – 12

3(NH4)CO3 N – 3; H – 12; C – 3; O – 9

4Al(C2H3O2)3 Al – 4; C – 24; H – 36; O – 24

2. What is the law of conservation of mass? During a chemical recation matter/mass cannot is neither created nor

destroyed

3. Balance the following chemical equations:

a. 1 KI + 4 F2 1 IF7 + 1 KF

b. 2 CuO + CuS 3 Cu + 1 SO2

c. 8 H3N + 3 Br2 6 NH4Br + 1 N2

d. 3 I2 + 10 HNO3 6 HIO3 + 10 NO + 2 H2O (HARD ONE)4. Correctly write out, then balance the following equations:

70

a. Nitrogen gas reacts with water to form hydrogen gas and trinitrogen pentaoxide

3 N2 + 10 H2O 10 H2 + 2 N3O5

b. Ammonium phosphide reacts with table salt to form ammonium chloride and sodium phosphide

1 (NH4)3P + 3 NaCl 3 NH4Cl + 1 Na3P

c. Vanadium III oxide reacts with calcium bromide to form vanadium III bromide and calcium oxide

1 V2O3 + 3 CaBr2 2 VBr3 + 3 CaO

d. Methanol (CH3OH) burns in a combustion reaction

2 CH3OH + 3O2 2 CO2 + 4 H2O

5. Identify the type of reaction in each of the following equations

a. CH2CH2 + Br CH3CH2Br SYNTHESIS

b. CaCO3 CaO + CO2 DECOMPOSITION

c. ZnF2 + Li LiF + Zn SINGLE DISPLACEMENT

d. C3H8 + O2 CO2 + H2O COMBUSTION

e. Ca(C2H3O2)2 + Na2CO3 CaCO3 + NaC2H3O2 DOUBLE DISPLACEMENT

71

72

Code Outcome description Understanding LevelLevel 1 Level 2 Level 3

3.1 Atomic structure Consistently identifies 3 of the following regarding subatomic particles: Names Charges Location in the atom Relative mass

Consistently identifies all of the items listed in level 1

3.2 Interpretation of the periodic table

Consistently demonstrates an understanding of 3 of the items listed below: Atomic mass Determining the number of neutrons

(including isotopes) Commonalities amongst groups and

periods Location of metals and non-metals Names of major groups

Consistently demonstrates an understanding of 4 of the items listed in level 1

Consistently demonstrates an understanding all of the items listed in level 1

3.3 Representing the Atom

Consistently demonstrates an understanding of 1 of the skills listed below: Bohr Model diagrams (protons,

neutrons, electron arrangement, shells) Lewis Dot Diagrams (nucleus, electron

arrangement)

Consistently demonstrates an understanding of all of the skills listed in level 1

3.4 Basic ionic bonding Consistently demonstrates an understanding of 2 of the skills listed below: Determine ionic charges from the

periodic table Determine the formula of a compound if

given the name Determine the name of a compound if

given the formula

Consistently demonstrates an understanding of all of the skills listed in level 1

3.5 Advanced ionic Consistently demonstrates an Consistently demonstrates an

bonding understanding of 2 of the skills listed below: Determine the formula of a compound

which includes a transition metal if given the name

Determine the name of a compound which includes a transition metal if given the formula

Determine the formula of a compound which includes a polyatomic ion if given the name

Determine the name of a compound which includes a polyatomic ion if given the formula

understanding of 3 or more of the skills listed in level 1

3.6 Covalent bonding Consistently demonstrates an understanding of 1 of the skills listed below: Determine the formula of a compound if

given the name Determine the name of a compound if

given the formula

Consistently demonstrates an understanding of 2 of the skills listed in level 1

3.7 Balancing chemical equations

Sometimes balances chemical equations (50-75%)

Usually balances chemical equations (75-90%)

Consistently balances chemical equations (90-100%)

3.8 Representing chemical equations

Consistently creates chemical equations from word equations (diatomic, reactants and products, transition metals, polyatomic ions)

3.9 Types of chemical reactions

Sometimes identifies types of chemical reactions (up to 60%)

Usually identifies types of chemical reactions (60-100%)