Science 1206 Unit 2: Chemistry Chemicals in Action Unit 2: Chemistry 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. 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. 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. Pure Substances A substance with constant composition throughout. Examples: Table salt (sodium chloride, NaCl), sugar (sucrose,C12H22O11), water (H2O), iron (Fe), copper and oxygen (O2). There are two types of pure substances - Element - Compound
(Cu), 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). 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). Read pgs. 172-174
Questions 1-3 pg.175 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. 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 + O2 ------> Fe2O3 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. 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)2 -----> KNO3 + PbI2 colourless reactants -----> coloured products
2. Solid (precipitate) produced potassium iodide + lead nitrate -----> potassium nitrate + lead iodide KI + Pb(NO3)2 -----> KNO3 + PbI2 -----> lead iodide will settle out 3. Gas produced zinc + hydrochloric acid -----> zinc chloride + hydrogen gas Zn + 2HCl -----> ZnCl2 + H2
-----> hydrogen gas bubbles out of solution 4. Energy change Hydrocarbon + oxygen -----> carbon dioxide + water vapour CH4 + O2 -----> CO2 + H2O -----> heat given off - energy given off (Exothermic)
- energy required (Endothermic) 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.
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.
Questions 4-8 pg.175 Chemicals and Safety (See pg. 176 & 658) Hazardous Household Product Symbols (HHPS) 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 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 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
Non-Metals Elements with the following properties Property lustre Property of Metal dull State - the states, at room
malleability brittle temperature, of all the elements are indicated through conductivity mostly insulators a text or colour code, depending on individual reactivity with acids no periodic table. state at room
solids, liquids, gases Periodic Table Sub Groups Periods horizontal rows are called periods represent the number of energy levels in atoms of an element Period 1 - 1 energy level Period 2 - 2 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
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 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) 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. 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. 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 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 23 - 11 = 12 neutrons) the number of electrons is equal to the number of protons (Ex. Sodium has 11 electrons)
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 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 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 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 23 - 11 = 12 neutrons
# electrons number of protons eg. Sodium atoms have 11 electrons Complete the following chart, for the first 18 elements, by listing the required information about individual atoms.
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.
Energy Level 1 2 3 Maximum # of Electrons 2 8
8 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. Electron Energy Level Diagrams - Atoms:
Chlorine Atom Nitrogen Atom Argon Atom Aluminum Atom Lithium Atom Beryllium
Read pgs. 184-187 Questions 1-3 pg. 187 Complete the chart on page 21 of you notes by filling in the missing information. 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. 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. Of course, this transfer of electrons will only occur if there are both metallic and non-metallic atoms in close proximity to each other.
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. 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. 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 Complete the charts and questions on pages 24-27 of your notes
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. 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 0 Read pgs. 184-187 Questions 4-8 pg. 187
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. 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. 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. 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 nonmetals (negative ions) the force of attraction exists between oppositely charged ions (opposites attract) dissolve in water and conduct electricity (electrolytes)
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) 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 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 Read pgs. 188-189 Questions 1,2 & 4 pg. 180/ 1 & 2 pg. 191 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. 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. 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. Remember that valence electrons are those found in the outermost energy level of an atom, or ion. 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 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
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 strontium and nitrogen aluminum and oxygen
magnesium and sulfur sodium and chlorine germanium and selenium zinc and phosphorous aluminum and arsenic
barium and phosphorous 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. 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 titanium (IV) & nitrogen vanadium and oxygen
nickel (III) and sulfur iron (III) and sulfur cobalt (II) and nitrogen lead (IV) and selenium bismuth (V) and astatine
gold (III) phosphorous 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. Polyatomic Formula Ion Name chlorite ClO2chromate CrO42- There are three complex
ions that do end in ide, namely cyanide, hydroxide and bisulfide, and these must be remembered. 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 Magnesium & chlorite Potassium & chromate
Nickel III & carbonate Strontium & cyanide Bismuth III & thiocyanate Calcium & nitrite Titanium III & permanganate
Type 3: Positive Complex Ion + Negative Complex Ion Ex. A compound formed by Ammonium and Carbonate Ammonium & iodate Ammonium & cyanide Ammonium & stearate
Ammonium & sulfate Ammonium & bisufite Ammonium & hydroxide Ammonium & carbonate Ammonium & benzoate
Type 4: Positive Complex Ion + Nommetallic Ion Ex. A compound formed by ammonium and oxygen. Ammonium & oxygen Ammonium & fluorine Ammonium & nitrogen
Ammonium & sulfur Ammonium & phosphorous Ammonium & chlorine Ammonium & tellurium Ammonium & iodine
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.
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. 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 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. Complete chart on pg. 45 Questions 1,3-6 pg. 195 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. 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 2times 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 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 1times 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
Complete chart on pg. 47 Questions 7-9 pg. 47 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. 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. Complete chart on pg. 48 Questions 1-4, 6 & 7 pg. 49 Complete charts on pgs. 50-52
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 CoCl 2 * 2H2O. 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. Numbe r
Prefix 1 2 3 4 5
6 7 8 9 10
mon o Di Tri Tetra Penta
Hexa Hept a Octa Nona Deca
Complete chart on pg. 53 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. 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. 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 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 Numbe r Prefix 1 2
3 4 5 6 7 8
9 10 mon o Di Tri
Tetra Penta Hexa Hept a Octa
Nona Deca 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
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
# Compound Name 1 Nitrogen monoxide 2 3
N2O Nitrogen dioxide 4 5 Compound Formula N2O3 Dinitrogen tetraoxide
Complete chart on pg. 56-57 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; #
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 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; # 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 Other unique Molecules Atoms of the elements phosphorous and sulfur will also combine with each other to form molecular compounds. # Compound Name
Compound Formula 1 Sulfur S8 2
Phosphorus P4 Monatomic Elements The noble gases are monatomic elements meaning they form molecular compounds as single atoms. # Compound Name
Compound Formula 1 Helium He 2
Neon Ne 3 Argon Ar 4
Krypton Kr 5 Xenon Xe
6 Radon Rn Complete chart on pg. 62 & 63 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. 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 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. 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. # Compound Formula Compound Name 1 HCl(g)
Hydrogen chloride 2 HCl(aq) Hydrochloric acid 3
H2SO4(g) Hydrogen sulfate 4 H2SO4(aq) Sulfuric acid 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 H2SO4(aq) and HNO2(aq) Molecule contains hydrogen, oxygen and another non-metal. These compounds form when hydrogen combines with polyatomic ions that contain oxygen
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. 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. Rule #1: Chemical Formulas: If the acid name begins with hydro and ends in ic, then the intermediate ionic name would have been hydrogen rootide. 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).
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.
Rule #2: Chemical Formulas: If the acid name ends with ic, then the intermediate ionic name would have been hydrogen rootate. 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). 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.
Rule #3: Chemical Formula: If the acid name ends with ous, then the intermediate ionic name would have been hydrogen rootite. 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). 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)