1. Science 1206
Unit 2: Chemistry

2. Chemicals in Action
Unit 2: Chemistry

3. CHEMISTRY TERMINOLOGY
Chemistry is the study of matter, its properties, and its changes or transformations. Matter is anything that has mass and takes up space. Matter can be categorized into two distinct types of materials. These are mixtures and pure substances.

4. Mixtures Two or more substances, combined in varying proportions - each retaining its own specific properties. The components of a mixture can be separated by physical means, i.e. without the making and breaking of chemical bonds. Examples: Air, table salt thoroughly dissolved in water, milk, wood, and concrete.

5. There are two types of mixtures
There are two types of mixtures. Heterogeneous Mixture A mixture in which the properties and composition are not uniform throughout the sample. A mixture with more than one visible substance. Examples: Pizza and concrete. Homogeneous Mixture Mixture in which properties and composition are uniform throughout the sample. A mixture with only one visible substance. Such mixtures are termed solutions. Examples: Air and table salt thoroughly dissolved in water.

6. Pure Substances A substance with constant composition throughout
Pure Substances A substance with constant composition throughout. Examples: Table salt (sodium chloride, NaCl), sugar (sucrose,C12H22O11), water (H2O), iron (Fe), copper (Cu), and oxygen (O2). There are two types of pure substances - Element - Compound

7. Element A substance that cannot be separated into two or more substances by ordinary chemical (or physical) means. We use the term ordinary chemical means to exclude nuclear reactions. Elements are composed of only one kind of atom. Examples: Iron (Fe), copper (Cu), and oxygen (O2).

8. Compound
A substance that contains two or more elements, in definite proportion by weight.
The composition of a pure compound will be invariant, regardless of the method of preparation. Compounds are composed of more than one kind of atom. The term molecule is often used for the smallest unit of a compound that still retains all of the properties of the compound.
Examples: Table salt (sodium chloride, NaCl), sugar (sucrose,C12H22O11), and water (H2O).

9. Read pgs. 172-174 Questions 1-3 pg.175

10. Chemicals and Chemical Change
Properties of Matter - matter has specific chemical and physical properties. Physical property a characteristic of a substance. A substance can have specific physical properties such as state, colour, texture, shape, mass, odour, density, taste, etc. Physical change a change in the form of a substance. For instance, from solid to liquid or liquid to gas or solid to gas, without changing the chemical composition of the substance. As we will see later, chemical bonds are not broken in a physical change. Examples: Boiling of water and the melting of ice.

11. Chemical property a characteristic behaviour that occurs when a substance changes to a new substance by reorganization of the atoms. Chemical change The change itself involves a chemical reaction. Example 1 Iron will react with oxygen to produce iron oxide(rust). Iron + oxygen > iron oxide Fe + O > Fe2O3

12. Example 2 Water will decompose, if induced by an electric current, to produce gaseous hydrogen and gaseous oxygen. Water > hydrogen + oxygen H2O > H2 + O2 The starting materials in chemical reactions are called reactants, and the new materials produced are called products.

13. Evidence of Chemical Change We use several pieces of evidence to verify if a chemical change has occurred. These include: 1. Colour changes potassium iodide + lead nitrate -----> potassium nitrate + lead iodide KI + Pb(NO3) > KNO3 + PbI2 colourless reactants -----> coloured products

14. 2. Solid (precipitate) produced potassium iodide + lead nitrate -----> potassium nitrate + lead iodide KI + Pb(NO3) > KNO3 + PbI > lead iodide will ‘settle out’ 3. Gas produced zinc + hydrochloric acid -----> zinc chloride + hydrogen gas Zn + 2HCl -----> ZnCl2 + H > hydrogen gas bubbles out of solution

15. 4. Energy change Hydrocarbon + oxygen -----> carbon dioxide + water vapour CH4 + O > CO2 + H2O -----> heat given off - energy given off (Exothermic) - energy required (Endothermic)

16. Chemical Tests (pg. 174)
Chemists can use chemical tests, or distinctive chemical reactions, to identify unknown gases or other substances. the presence of oxygen gas, for example, is tested using a glowing splint. The glowing splint ‘bursts into flames’ when placed in a container that has concentrated oxygen gas present.

17. the presence of hydrogen gas, for example, is tested using a flaming splint. The flaming splint causes the gas to ‘pop’ when placed in a container that has concentrated hydrogen gas present. the presence of carbon dioxide gas, for example, is tested using a limewater solution. The carbon dioxide gas is bubbled into the limewater solution. If the solution turns milky the gas is carbon dioxide.

18. Questions 4-8 pg.175

19. Chemicals and Safety (See pg. 176 & 658)
Hazardous Household Product Symbols (HHPS)

20. Elements and the Periodic Table
The Periodic Table a structured arrangement of elements that helps us to explain and predict physical and chemical properties. the table was designed by Dmitri Mendeleev represents on of the most important advancements in science

21. The periodic table is generally arranged
with metals towards the left side of the table (staircase).
nonmetals on the right side of the table (staircase).
hydrogen the exception - found in the top left corner but behaves mostly as a nonmetal

22. Metals Elements with the following properties Property
Property of Metal
lustre
shiny
malleability
malleable
conductivity
conductors
reactivity with acids
yes
state at room
temperature solids, except mercury
Elements with the following properties

23. Non-Metals Elements with the following properties
State - the states, at room temperature, of all the elements are indicated through a text or colour code, depending on individual periodic table.
Property
Property of Metal
lustre
dull
malleability
brittle
conductivity
mostly insulators
reactivity with acids no
state at room
solids, liquids, gases

24. Periodic Table Sub Groups
Periods horizontal rows are called periods represent the number of energy levels in atoms of an element Period energy level Period energy levels and so on Families groups of elements in the same vertical column elements with similar physical and chemical properties some of these groups include

25. Family Sub Groups
Alkali Metals also called Group IA elements far left column of the periodic table very reactive form compounds that are mostly white solids and are very soluble in water all shiny, silvery metals Alkaline Earth also called Group IIA elements include magnesium (Mg) and calcium (Ca) shiny, silvery metals form compounds often insoluble in water

26. Noble Gases far right column of the periodic table
also called Group VIIIA elements
include helium (He) and neon (Ne)
inert - generally do not form compounds
Halogens second column from the right
nonmetallic elements
also called Group VIIA elements
all poisonous elements
react readily with alkali metals
include chlorine (Cl), bromine (Br)

27. Hydrogen. a very special element. It is located in Group IA and
Hydrogen a very special element. It is located in Group IA and Group VIIA. It sometimes behaves as a metal, but behaves mostly like a nonmetal.

28. Elements and Atomic Structure
Atoms are the basic building blocks of all forms of matter in our known universe. Everything we see and touch is made up of atoms. An atom is the smallest particle of an element that cannot be further broken down by physical or chemical means. The atom also has a very particular structure.

29. Bohr-Rutherford Model of the Atom There are many models of the atom that have been proposed throughout history. The Bohr- Rutherford Model of the atom suggests there are three sub-atomic particles in the atom. These are located in two regions of the atom, namely - the nucleus - the electron cloud

30. Atom neutral in charge # of protons equal to the atomic number (ex
Atom neutral in charge # of protons equal to the atomic number (ex. Sodium has 11 protons) # of neutrons equals the atomic # - atomic mass # (eg. Sodium has = 12 neutrons) the number of electrons is equal to the number of protons (Ex. Sodium has 11 electrons)

31. Nucleus dense positively charged core of the atom composed of protons and neutrons. protons are relatively small in volume (size) protons are heavy (high mass) particles protons are positively charged particles neutrons are relatively small in size neutrons are heavy (high mass) particles neutrons are neutral particles the nucleus has an overall net positive charge the nucleus makes up the vast majority of the mass of the atom

32. Electron cloud region surrounding the nucleus where electrons are found electrons have relatively large volumes(size) electrons have very little mass electrons are negatively charged particles electrons are located in energy levels first energy level two electrons second energy level eight electrons third energy level eight electrons

33. the electron cloud has an overall net negative charge makes up the vast majority of the volume of the atom the electrons spin in shells around the nucleus of the atom

34. Sub-atomic Particles The number of protons, neutrons and electrons found in an atom can be determined using the periodic table.
# protons
atomic # from periodic table
eg. Sodium(Na) atoms have 11 protons
# neutrons
atomic mass # - atomic #
eg. Sodium atoms have = 12 neutrons
# electrons
number of protons
eg. Sodium atoms have 11 electrons

35. Complete the following chart, for the first 18 elements, by listing the required information about individual atoms.

36. Electron Energy Level Diagrams (Bohr Diagrams) for Atoms
These diagrams indicate the main features of the atom of an element. Bohr diagrams are used to represent the number of protons and neutrons found in the nucleus of an atom and the number of electrons in the electron cloud. They also represent the arrangement of electrons in various orbits, or energy levels, around the nucleus. Each electron energy level has a specific and definite number of electrons.

37. Energy Level
Maximum # of Electrons 1 2 8 3

38. The outer-most energy level that contains electrons is called the valence energy level, or valence shell and electrons located here are called valence electrons. The Bohr diagram also show the name and symbol of the atom. As well, atoms are electrically neutral since the number of positively charged protons is equal to the number of negatively charged electrons.

39. Electron Energy Level Diagrams - Atoms:
Chlorine Atom
Nitrogen Atom
Argon Atom

40. Aluminum Atom
Lithium Atom
Beryllium

41. Read pgs. 184-187 Questions 1-3 pg. 187

42. Complete the chart on page 21 of you notes by filling in the missing information.

43. Compounds and Bonding
Atoms to Ions It is important to note that atoms tend to be unstable structures. They are generally unstable because the outermost energy levels that contain electrons are not filled. Atoms want to gain or lose electrons to ensure that their outer energy levels are filled with electrons, and to ensure they are stable structures.. However, when atoms gain or lose electrons they become ions.

44. Atoms become ions to attain an electron energy level arrangement like the nearest noble gas. The noble gases have stable atoms because their outermost energy levels are filled with electrons. In order to attain this stable electron structure metallic atoms, which have weak attractions for valence electrons, will lose electrons to become like their nearest noble gas. This will result in the metallic ions having a positive charge. The non-metallic atoms, which have strong attractions for valence electrons, will gain electrons to become stable like their nearest noble gas. This will result in the non-metallic ions having a negative charge.

45. Of course, this transfer of electrons will only occur if there are both metallic and non-metallic atoms in close proximity to each other.

46. Electron Energy Level Diagrams for Ions:
Aluminum Ion
Aluminum will give up three electrons to become an ion, and to be stable like its nearest noble gas - neon
Because aluminum loses three electrons to become an ion it attains an overall charge of 3+.
Note that the symbol changes, as well, to represent the charge of the ion.

47. Electron Energy Level Diagrams for Ions:
Oxygen Ion
Oxygen will gain two electrons to become an ion, and to be stable like its nearest noble gas - neon.
Because oxygen gains two electrons to become an ion it attains an overall net charge of 2-.
When non-metallic atoms gain, or lose, electrons to become ions their name changes. Oxygen atom, for example, changes to become oxide ion.

48. Electron Energy Level Diagrams for Ions:
Beryllium Ion
Beryllium will give up two electrons to become an ion and to be stable like its nearest noble gas - helium.
Because beryllium loses two electrons to become an ion it attains

49. Complete the charts and questions on pages 24-27 of your notes

50. Remember that families of elements have similar chemical and physical properties. These families of elements will gain, or lose, specific numbers of electrons to attain a stable ‘noble gas like’ electron arrangement. All elements in group IA, for example will lose one electron to be like the nearest noble gas. The other families are as follows.

51. Group
Gain or Lose e- to become ion
Ion Charge
Group IA (alkali metals)
lose of one electron to become an ion +1
Group IIA (alkaline earth)
lose two electrons to become an ion +2
Group IIIA
lose three electrons to become an ion +3
Group VA
gain three electrons to become an ion -3
Group VIIA (halogens)
gain two electrons to become an ion -1
Group VIIIA (noble gases)
gain one electron to become an ion

52. Read pgs. 184-187 Questions 4-8 pg. 187

53. Bonding: How Elements Form Compounds
Bonding occurs between elements when conditions are such that there are two, or more, elements present in close proximity to each other. A force of attraction between these elements results in the formation of a bond, and a compound is created. There are many different types of bonding that can occur between atoms of elements.

54. In Science 1206, we focus on two main types, namely; 1. Ionic bonding 2. Covalent (molecular) bonding The type of bonding that will actually occur depends on the types of atoms that are in close proximity to each other. In order for these atoms to create a bond there must be a force of attraction between the them.

55. In order for sodium chloride to be formed there must be sodium and chlorine atoms in close proximity to each other, and there must be a force of attraction between them. This force of attraction will create sodium chloride, an ionic compound.

56. Ionic Compounds
made up of negative and positive ions that have resulted from the transfer of electrons from a metal to a nonmetal. electrons are transferred from metals (positive ions) to non-metals (negative ions) the force of attraction exists between oppositely charged ions (opposites attract) dissolve in water and conduct electricity (electrolytes)

57. Molecular Compounds
formed when non-metals combine with non-metals a force of attraction exists between the positive nucleus of one atom and the negatively charged electrons of the other atom(s). electrons are shared between the two non-metals do not conduct electricity in solution (non-electrolytes)

58. Chemical Formula: Ratios of Atoms
The compounds that are created will have specific ratios of one atom(element) to another. This ratio is represented using a chemical formula. This chemical formula contains a combination of symbols that represent a particular compound and denotes the relative numbers of each element in the compound. i.e.: NaCl This chemical formula indicates that there is one sodium for every one chlorine ion(chloride) in the compound

59. i.e.: CaCl2 This chemical formula indicates that there is one calcium ion for every two chlorine ions(chloride) in the compound I.e.: P4O10 This chemical formula indicates that there are four phosphorous ions for every ten oxygen ions(oxide) in the compound

60. Read pgs. 188-189 Questions 1,2 & 4 pg. 180/ 1 & 2 pg. 191

61. Ionic Bonding and Ionic Compounds
Ionic bonding occurs when atoms(ions) with opposite charges are in close proximity to each other. Atoms of elements attain charges by gaining, or losing, electrons to form ions. The atoms will gain or lose electrons if there are elements present that have different attractions for electrons.

62. Metallic atoms, for example, have weak attractions for electrons while non-metallic atoms have strong attractions for electrons. As a result, when metallic and non-metallic atoms are in close proximity to each other the non-metallic atoms steal the some of the electrons away from the metallic atom. This results in a positively charged metallic ion, and a negatively charged non-metallic ion. The number of valence electrons gained, or lost, will depend on the atoms present and on the number of electrons they need to transfer to obtain a noble gas like electron structure - to become stable ions.

63. Sodium, for example has a weak attraction for electrons while oxygen has a strong attraction for electrons. When these elements are in close proximity to each other there will be a transfer of electrons from the sodium to the oxygen. The sodium atoms will end up with a positive charge and the oxygen will end up with a negative charge. These ions, with opposite charges, are then attracted to each other(opposites attract), creating the compound Na2O. Opposites attract in ionic bonding.

64. Remember that valence electrons are those found in the outermost energy level of an atom, or ion.

65. Writing Chemical Formulas for Ionic Compounds
There are four distinct combinations of ions that determine the formation of an ionic compound. These criteria are dependent on the types of atoms, or ions, that are present to form the compound. These combinations of ions are: Type 1. metal ion + non-metal ion Type 2. metal ion + negatively charged complex ion(polyatomic ion) Type 3. positively charged complex ion + negatively charged complex ion Type 4. positively charged complex ion + non-metal

66. Rules for Determining Chemical Formula of Ionic Compounds.
Rule 1 Write the symbols, with the cation(+) first Rule 2 Write the Ionic charge above each symbol to indicate the stable ion each element forms Rule 3 Choose the number of ions to balance each charge. Use the cross-hatch method. Let the number that represents the ionic charge be the subscript for the oppositely charged ion. This creates a neutrally charged compound. Rule 4 Write the formula using subscripts

67. Type 1: Metal Ion(cation) + Nonmetal Ion(anion)
Ex. A compound formed by Aluminum and Bromine - note the overall charge on the compound - do several more examples Complete the chart on page 33-39

68. strontium and nitrogen
aluminum and oxygen
magnesium and sulfur

69. germanium and selenium
sodium and chlorine
germanium and selenium
zinc and phosphorous

70. barium and phosphorous
aluminum and arsenic
barium and phosphorous

71. Type 1: (continued) Multivalent Metal Ion(cation) + Nonmetal Ion(anion)
Note: Some metallic atoms are capable of forming ions with more than one possible ionic charge. These elements are referred to as multivalent. If the charge of the multivalent ion is not given then you must always choose the top one as this is the one that occurs most often.

72. Ex. A compound formed by Iron (II) and Nitrate - note the overall charge on the compound - do several more examples Complete the chart on page 39-40

73. titanium (IV) & nitrogen vanadium and oxygen nickel (III) and sulfur

74. cobalt (II) and nitrogen
iron (III) and sulfur
cobalt (II) and nitrogen
lead (IV) and selenium

75. bismuth (V) and astatine gold (III) phosphorous

76. Type 2: Metal(cation) + Negative Complex(Polyatomic) Ion(anion)
The complex ions are a special group of ions that have formed when a group of atoms have combined to produce a single, stable complex ion with an overall charge. The elements boron and oxygen can form a single stable ion called borate. This complex ion has an over charge of 3- and will have the formula BO3-. It is also located in the table of complex ions. Note the names of the complex ions ending with the suffixes ate and ite.

77. Polyatomic Ion Name
Formula
chlorite
ClO2-
chromate
CrO42-
There are three complex ions that do end in ide, namely cyanide, hydroxide and bisulfide, and these must be remembered.

78. As well, it is important to note that if there is more than one complex ion in the chemical formula of the compound, brackets must be used to obtain the proper ratio for the complex ion.
Ex. A compound formed by Indium and Nitrate

79. Magnesium & chlorite
Potassium & chromate
Nickel III & carbonate

80. Bismuth III & thiocyanate
Strontium & cyanide
Bismuth III & thiocyanate
Calcium & nitrite

81. Titanium III & permanganate

82. Type 3: Positive Complex Ion + Negative Complex Ion
Ex. A compound formed by Ammonium and Carbonate

83. Ammonium & iodate
Ammonium & cyanide
Ammonium & stearate

84. Ammonium & sulfate
Ammonium & bisufite
Ammonium & hydroxide

85. Ammonium & carbonate
Ammonium & benzoate

86. Type 4: Positive Complex Ion + Nom-metallic Ion
Ex. A compound formed by ammonium and oxygen.

87. Ammonium & oxygen
Ammonium & fluorine
Ammonium & nitrogen

88. Ammonium & phosphorous
Ammonium & sulfur
Ammonium & phosphorous
Ammonium & chlorine

89. Ammonium & tellurium
Ammonium & iodine

90. Rules for Writing Compounds Names of Ionic Compounds.
Rule 1: - Name the positive ion(cation) first. - If the ion is milti-valent then use roman numerals to indicate the charge on the ion. Rule 2: - Name the anion second. - If the anion is a single element (non-metal) remove the suffix and append "-ide". - If the anion is a complex ion write as given in the table of complex ions.

91. Example 1: A compound formed by aluminum and bromine The cation is aluminum so it gets written in full. The anion is bromine so we drop the ine and add ide. Bromine becomes bromide. Thus the compound name is aluminum bromide.

92. Example 2: A compound formed by magnesium and oxygen
The cation is magnesium so it gets written in full.
The anion is oxygen so we drop the ygen and add ide. Oxygen becomes oxide.
Thus the compound name is magnesium oxide

93. Example 3: A compound is formed from copper and chlorine
Example 3: A compound is formed from copper and chlorine. Note that copper is multivalent and if you do not know the charge on the ion then choose the ion at the top. In this case it is Cu2+. The name of the compound must now indicate this charge. The cation is copper so it gets written in full, with the charge indicated in roman numerals - copper (II) The anion is chlorine so we drop the ine and add ide. Chlorine becomes chloride Thus the compound name is copper (II) chloride.

94. Complete chart on pg. 45 Questions 1,3-6 pg. 195

95. Multivalent Ion Note:. If you are given the chemical formula for the
Multivalent Ion Note: If you are given the chemical formula for the compound then the charge on the multivalent ion can be determined by multiplying the charge on the negatively charged anion with the subscript denoting the ratio of the anion to the cation. Use the cross method to determine the charge on the multivalent ion.

96. Example 4: TiO2 - the titanium is multivalent
In this case, you can determine the charge on the titanium by multiplying the charge on the anion by the ratio number of the ion 2- times 2 = 4-.
Since the compound must be neutral then there must be a net positive charge of 4+, and since there is only one titanium ion it must have a charge of 4+.
Name: Titanium (IV) oxide

97. Example 5: SnCl2 - the tin is multivalent
In this case, you can determine the charge on the tin by multiplying the charge on the anion by the ratio number of the ion 1- times 2 = 2 -.
Since the compound must be neutral then there must be a net positive charge of 2+, and since there is only one tin ion it must have a charge of 4+.
Name: Tin(II) chloride

98. Complete chart on pg. 47 Questions 7-9 pg. 47

99. Example 6: A compound is formed from calcium and sulfate
Example 6: A compound is formed from calcium and sulfate. Note name ending of sulfate as ‘ate’. This indicates this ion is a complex ion. Ions ending in ‘ite’ are also complex ions. The cation is calcium so it gets written in full. The anion is sulfate, a complex ion, so it gets written as in the table of complex ions. Thus the compound name is calcium sulfate.

100. Example 7: A compound is formed from aluminum and sulfite
Example 7: A compound is formed from aluminum and sulfite. Note name ending of sulfite as ‘ite’. This indicates this ion is a complex ion. Ions ending in ‘ate’ are also complex ions. The cation is aluminum so it gets written in full. The anion is sulfite, a complex ion, so it gets written as in the table of complex ions. Thus the compound name is aluminum sulfite.

101. Complete chart on pg. 48 Questions 1-4, 6 & 7 pg. 49

102. Complete charts on pgs. 50-52

103. Hydrates
Hydrated substances are ionic compounds that contain water as part of their crystalline structure. Cobalt chloride, for example, can incorporate a water component depending on how the compound is formed. The hydrated ionic compound produced is called cobalt (II) chloride dihydrate. The ‘di’ prefix is used to indicate the amount, or number, of water molecules contained in the compound. The chemical formula for this ionic compound will be CoCl2 * 2H2O.

104. When writing the compound name for a hydrated compound it is necessary to indicate the amount, or number, of water molecules contained in the compound by using the a prefix system.
Example:
NaCl * 3H2O Note that there are three water molecules that are contained in this compound. The prefix for three is ‘tri’, so the name of this compound will be sodium chloride trihydrate.

105. Number 1 2 3 4 5 6 7 8 9 10
Prefix
mono Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca

106. Complete chart on pg. 53

107. Molecular Compounds
Many compounds are formed when non-metallic atoms are linked together to form electrically neutral compounds called molecules. These compounds are formed when one or more non-metallic elements combine through a process called covalent bonding. Molecular compounds share electrons to become stable, like its nearest noble gas, and to attain an overall neutral electric charge.

108. In molecules, the electrons of one atom are attracted to the nucleus of its neighbouring atom, and vive-versa. Carbon dioxide (CO2), for example, is a molecular compound since both carbon and oxygen are non- metallic elements. Carbon tetrafluoride is another molecular compound. Its formula will be CCl4. These compounds may contain just two atoms of the same element (diatomic molecules), such as hydrogen gas (H2) and oxygen gas (O2). Most molecules are complex and will contain many more atoms such as sugar (C12H22O11). A compound in blood, for example, has the chemical formula C3032H4816O872N780S8F4.

109. Molecular Bonding( Covalent bonds)
bonds in which electrons are shared between the nuclei of atoms.
molecules are electrically neutral
covalent bonds form to complete the outer energy level of atoms and allow the atoms to have a noble gas like electron structure

110. Rules for Naming Molecular Compounds
The first element in the formula is written with its full name.
The second element is written as if it were an ion using “ide”
Prefixes are used to denote the numbers of atoms present.
The prefix mono- is only used to name the second element.
For example, CO is carbon monoxide, not monocarbon monoxide. Sometimes, however, the prefix mono is not used in naming the second element present. If there is no prefix the assumption is that mono should be there.
Common names are preferred where compounds have common names

111. Number 1 2 3 4 5 6 7 8 9 10
Prefix
mono Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca

112. Example 1: P4O10 Note the prefixes for both the phosphorous and oxygen atoms. The prefix 4 is tetra and 10 is deca. These will now be used to name the compound. As well, the name of the second atom must be changed to end in ‘ide’. Name: tetraphosphorous decaoxide

113. Example 2: CF4
Note the prefixes for both the carbon and fluorine atoms. The prefix 1 is mono and 4 is tetra. These will now be used to name the compound. Once again, the name of the second atom must be changed to end in ‘ide’. Remember that the rules indicate the prefix ‘mono’ does not get used with the first atom.
Name: Monocarbon tetrafluoride becomes carbon tetrafluoride

114. Dinitrogen tetraoxide#
Compound Name
Compound Formula 1
Nitrogen monoxide 2
N2O 3
Nitrogen dioxide 4
N2O3 5
Dinitrogen tetraoxide

115. Complete chart on pg

116. Compounds with Common Names.
There is a large group of molecular compounds that have been used over a long period of time and, as a result, have common names. These common names are the ones that must be used when writing the compound name for these substances. These compounds include the following;

117. #
Compound Name
Compound Formula 1
Water
H20 2
Ozone O3 3
Ammonia
NH3 4
Methane
CH4 5
Sucrose
C12H22O11 6
Methanol
CH3OH 7
Ethanol
C2H5OH 8
Hydrogen peroxide
H2O2

118. Diatomic Molecules
Other non-metallic atoms are capable of forming compounds by sharing electrons with other atoms of the same element. These molecular compounds are referred to as diatomic molecules, and they exist in nature as bonded pairs of the same element. The diatomic molecules are;

119. #
Compound Name
Compound Formula 1
Hydrogen H2 2
Nitrogen N2 3
Oxygen O2 4
Fluorine F2 5
Chlorine
Cl2 6
Bromine
Br2 7
Iodine I2 8
astatine
At2

120. Other unique Molecules
Atoms of the elements phosphorous and sulfur will also combine with each other to form molecular compounds.

121. #
Compound Name
Compound Formula 1
Sulfur S8 2
Phosphorus P4

122. Monatomic Elements
The noble gases are monatomic elements meaning they form molecular compounds as single atoms.

123. #
Compound Name
Compound Formula 1
Helium He 2
Neon Ne 3
Argon Ar 4
Krypton Kr 5
Xenon Xe 6
Radon Rn

124. Complete chart on pg. 62 & 63

125. Hydrogen Compounds
Hydrogen compounds are a special group of molecular compounds that are composed of a hydrogen atom combined with another non- metal atom or with a negatively charged complex ion. Hydrogen compounds are often said to have an ‘identity crisis’. They are molecular, covalently bonded compounds that behave as ionic compounds.

126. Hydrogen compounds - may be solid, liquid or gas at room temperature
Hydrogen compounds - may be solid, liquid or gas at room temperature. - are soluble in water - form coloured and colourless solutions - form solutions that conduct electricity (electrolytes) - form solutions which turn blue litmus red

127. Most hydrogen compounds are named as acids
Most hydrogen compounds are named as acids. A couple of exceptions include HCl(g), H2S(g) and HCN(g). These are named as if they were ionic.

128. Acids
When hydrogen compounds are placed in water they ionize (break apart into individual ions) to become acids. Because the acids are ionized in water they are given the state designation aqueous(aq) - meaning dissolved in water. All acids must have this (aq) designation.

129. #
Compound Formula
Compound Name 1
HCl(g)
Hydrogen chloride 2
HCl(aq)
Hydrochloric acid 3
H2SO4(g)
Hydrogen sulfate 4
H2SO4(aq)
Sulfuric acid

130. Any compound whose chemical formula begins with a hydrogen cation (H+) is an acid . There are two different types of acids. They are;
Type
Explanation
Example
Binary
Molecule contains hydrogen and one other non-metal
HF(aq) and HCl(aq)
OXO
Molecule contains hydrogen, oxygen and another non-metal.
These compounds form when hydrogen combines with polyatomic ions that contain oxygen
H2SO4(aq) and HNO2(aq)

131. Rules for Naming, and Writing Chemical Formulas, for Acids
Because of the nature of acids, being molecular compounds that behave as if they were ionic compounds, a special naming system was developed for these molecules. A special chart, called Naming Acids, is present in the periodic table to help accommodate this unique naming system. The general rule is to create an intermediate ionic name to represent the molecules chemical formula or compound name. This intermediate ionic name is then used in conjunction with the naming acids table in the periodic table.

132. Rule #1: Naming:
Acids with intermediate ionic names ending with ‘ide’. The root of the anion name is combined with ‘hydro’ and ‘ic’ to create the name of the acid. HCl(aq) has the intermediate ionic name hydrogen chloride - ending in ide. As a result the root of the anion chloride is combined with hydro and ic to produce the acid name hydrochloric acid. HCN(aq) has the intermediate ionic name hydrogen cyanide - ending in ide. As a result the root of the anion cyanide is combined with hydro and ic to produce the acid name hydrocyanic acid.

133. Rule #1: Chemical Formulas:
If the acid name begins with ‘hydro’ and ends in ‘ic’, then the intermediate ionic name would have been hydrogen ‘root’ide. The ions are then balanced as if they were ionic. Hydrobromic acid, for example, has an acid name beginning with ‘hydro’ and ending in ‘ic’. Its intermediate ionic name will be hydrogen bromide. The ions present will be H+ and Br -. Using the cross-hatch we can predict the formula to be HBr(aq).

134. Rule #2: Naming:
Acids with intermediate ionic names ending with ‘ate’. The root of the anion name is combined with ic to create the name of the acid. H2CO3(aq) has the intermediate ionic name hydrogen carbonate - ending in ‘ate’. As a result the root of the anion carbonate is combined with ‘ic’ to produce the acid name carbonic acid. H2SO4(aq) has the intermediate ionic name hydrogen sulfate - ending in ‘ate’. As a result the root of the anion sulfate is combined with ‘ic’ to produce the acid name sulfic acid. The anions sulfur and phosphorous retain their ur and or designation when named as acids. The acid name then is sulfuric acid.

135. Rule #2: Chemical Formulas:
If the acid name ends with ‘ic’, then the intermediate ionic name would have been hydrogen ‘root’ate. The ions are then balanced as if they were ionic. Chloric acid, for example, has an acid name ending in ‘ic’. Its intermediate ionic name will be hydrogen chlorate. The ions present will be H+ and ClO3-. Using the cross-hatch we can predict the formula to be HClO3(aq).

136. Rule #3: Naming:
Acids with intermediate ionic names ending with ‘ite’. The root of the anion name is combined with ‘ous’ to create the name of the acid. HNO2(aq) has the intermediate ionic name hydrogen nitrite - ending in ite. As a result the root of the anion nitrite is combined with ‘ous’ to produce the acid name nitrous acid. H2SO3(aq) has the intermediate ionic name hydrogen sulfite - ending in ite. As a result the root of the anion sulfite is combined with ‘ous’ to produce the acid name sulfous acid. The anions sulfur and phosphorous retain their ‘ur‘ and ‘or’ designation when named as acids. The acid name then is sulfurous acid.

137. Rule #3: Chemical Formula:
If the acid name ends with ‘ous’, then the intermediate ionic name would have been hydrogen ‘root’ite. The ions are then balanced as if they were ionic. Sulfurous acid, for example, has an acid name ending in ‘ous’. Its intermediate ionic name will be hydrogen sulfite. The ions present will be H+ and SO3 2-. Using the cross-hatch we can predict the formula to be H2SO3(aq).

138. Acids with COO- groups. Any acid with a COO- group at the end, or beginning, of the its chemical formula will need to have the Hydrogen(H+ ) cation placed next to the COO- group. Acetic acid, for example, is composed of the cation H+ and the anion acetate CH3COO-. These two ions will combine to create an acid with the chemical formula the HCH3COO(aq). Note, however, that the anion has a COO- group, and recall that the H+ ion has to be placed next to this COO- polyatomic ion group. This will result in the chemical formula, for this compound, being CH3COOH(aq)