Download Unit 1 – Physical Science and Chemical Reactions

UNIT 1 – PHYSICAL SCIENCE AND CHEMICAL REACTIONS
CHAPTER 5: CHEMICALS IN ACTION
chemistry: the study of matter
matter: anything that has mass and occupies space
MATTER
mixtures
SUBSTANCES
PURE
COMPOUNDS
ELEMENTS
molecular
ionic
metal
non-metal
pure substance: is one in which all the particles that make up the substance are the
same; has constant properties. Atoms are in a fixed ratio.
eg. (H2O(ℓ))
mixture: contains two or more pure substances not in a fixed ratio
 eg. sugar + water  sweet water
(solute) + (solvent)  (solution)
element: a pure substance composed of only one type of atom (eg. Iron (Fe))
 Cannot be broken down by ordinary chemical means (eg. heating, electricity,
filtration, distillation, magnets)
Metals
Non- metals
 Found on left hand side of staircase line
of P.T.
 Shiny
Dull
 Ductile, malleable
Brittle
R.H.S.
 Hard
soft
Ductile – the ability to be formed into a wire
Malleable – the ability to be stretched by hammering or rolling
compound: pure substance composed of at least two types of elements (or atoms)
 eg. water (H2O(ℓ)); salt (NaCl(s))
Ionic
Molecular
Composed of 2 or more ions ( metal + non-metal)
2 or more non-metallic
atoms
Held together by an ionic bond
covalent bond
(e.g.) NaCl- sodium chloride
CO2 – carbon dioxide
The Modern Periodic Table
Facts about hydrogen
Group IA
 1 valence electron
 Loses 1 electron to form a hydrogen ion
Group VIIA
 1 less electron than a Noble Gas
 A gas
Can gain 1 electron to form a hydride ion
Periodic Table
 Demitri Mendeleev created the first useable periodic table
 It became the most predictive device in all chemistry
 Elements are arranged in increasing order of Atomic Number (AN)
 The periodic table is divided into vertical columns (families/groups) and horizontal rows (periods)
Atomic Number: (AN) the number of protons (p+) in the nucleus of an atom
Group A Elements
IA ~ Alkali Metals
 Soft, reactive
IIA ~ Alkaline Earth Metals
 When compared to IA; harder or not as reactive
IIIA ~ Aluminum Group
IVA-VIA ~ Name according to first element in group
VIIA ~ Halogens
 Very reactive
VIIIA ~ Noble (inert) Gases
 All gases
 Non-reactive
Group B Elements (Transition Elements)
 When compared to Group A elements; harder and more colourful, with higher boiling and melting
points
metalloids: located on both sides of the staircase line and have characteristics of metals and nonmetals
Matter has a well defined underlying structure composed of the following:
1. Atoms
 The smallest whole part of an element that is still representative of the element
 A neutral particle composed of a nucleus containing protons (p+) and neutrons, and electrons
(e-) outside the nucleus (the number of electrons equals the number of protons)
2. Ions
 A charged atom (The charge is caused by the transfer of an electron (or electrons) from one
atom to another)
 The metallic atom (which loses an electron) has a positive charge and is called a cation
- eg. Na+, Al3+
 The non-metallic atom (which gains an electron) has a negative charge and is called an anion
- eg. Clˉ, N3ˉ
3. Molecules
 Occurs when non-metallic atoms share electrons (which then bond together into a group)
- eg. Methane (CH4)
H

. .
This is a Lewis Dot Diagram.
represent molecules.
C : H : C
It is used to
. .
H
Development of Atomic Theory
1. Dalton “Billiard Ball” Model (1803)
 States that the atom is composed of uniform matter, not divisible (cannot be broken down
into smaller parts)
2. Thompson “Raisin Bun” Model (1897)
 Evidence from experiments (electricity through gases in a vacuum tube) led Thompson to
believe that negative (-ve) charged particles were scattered throughout the positive atom like
raisins in a bun
 He called these negatively charged particles electrons. (He was the first)
3. Rutherford’s “Nuclear” Model (1914)
 By releasing (+ve) alpha particles at a thin gold foil Rutherford was astonished that a small
percentage were deflected at large angles. He suggested that the atom must have a small,
dense, positively charged inner core called a nucleus.
 He called these positively charged particles inside the nucleus protons. (He was the first)
 Outside the nucleus was a large area of mostly empty space containing electrons
 This work led one of Rutherford’s students (H. G. Mosely) to discover that the positive
charge inside the nucleus increased by one from element to element in Mendeleev’s periodic
table. From this he coined the phrase atomic number (the number of protons inside the nucleus
of an atom).
 This discovery led to modifications of Mendeleev’s periodic table – instead of listing the
elements in order of increasing atomic mass it would now be in order of increasing atomic number
 The mass of the protons was too small to account for the total atomic mass of the atom so
Rutherford predicted that there must be a neutral particle in the nucleus similar in mass to the
proton
 In 1932 James Chadwick demonstrated this particle, which he called a neutron. (He was the
first to name this particle)
4. Bohr “ Orbit” Model (1921)
 Electrons surround the nucleus of an atom in circular orbits, with each electron having a fixed
quantity of energy
 Electrons cannot exist between orbits but can move to unfilled orbits
 The higher the energy level, the further from the nucleus the electron is
 The maximum numbers of electrons in the first three orbits are 2, 8, and 8 respectively
(Octet Rule)
 An atom with the maximum number of electrons in the outermost energy level is stable
(unreactive)
ep+
e-
5. Quantum Mechanical Model (1920’s)
 Similar to Bohr’s model, except that the electrons are said to be found in orbitals (not orbits)
(an orbital is an area of space where an electron is likely to be found)
p+
 Other main features include:
- Atomic Number = Number of protons
- Number of protons = Number of electrons
- An energy level represents a specific value of energy of an electron and corresponds to a
general location
- Period number = Number of energy levels occupied by electrons
- The first three energy levels will have 2, 8, and 8 structures of electrons
NOTE: 2, 8, and 8 should be written as follows
started
-8e--
-8e-- -8e--
-2e-- -2e-- -2e-- Each lower energy level must be filled to its maximum before the next level is
- The electrons in the highest energy level are called valence electrons
- Group number = Number of valence electrons (one exception is helium) (Group A only)
- Max. number of electrons in each energy level = Max. number of atoms in each period
Energy Level Diagrams For Atoms
valence level: the outer-most energy level of an energy level diagram
NOTE: It is important to include atom when writing the name
 ex. Lithium atom (Students will be expected to format all energy level diagrams for atoms as
follows)
Valence Electrons (Valence Level):
- 1e-- 2e-Atomic Number:
3p+
Name: Lithium atom
Symbol:
Li
Energy Level Diagrams For Ions
ion: a charged atom
 Formed when a metal loses electrons to form a positive ion (cation) or when non-metals gain
electrons to form a negative ion (anion)
 All atoms gain/lose electrons to become like the nearest noble gas (become stable)
When naming ions:
 Metal (cation) names: stay the same
 Non-metal (anion) names: change (last three letters become “ide”)
ex. Lithium ion (Students will be expected to format all energy level diagrams for ions as follows)
Valence Electrons (Valence Level):
- 2e-3p+
Atomic Number:
Name:
Symbol:
Li+
Lithium ion
 ex. Fluoride ion (The atom Fluorine becomes the ion Fluoride)
- 8e--
Valence Electrons (Valence Level):
- 2e-Atomic Number:
9p+
Name: Fluoride ion
Fˉ
Symbol:
Isotopes
 Atomic number = Number of protons in the nucleus of an atom
 Atomic mass (mass number) = Sum of the protons and neutrons in the nucleus of an atom
isotopes: forms of an element that have the same number of protons but different numbers of
neutrons
Two Ways to Write Isotopes
 ex. Argon
Atomic mass (AM)
 ex. Argon
Atomic number (AN)
argon-40
40
18
Ar
Atomic mass
Retrieving Information About Isotopes
 Atomic mass is given in both naming forms
 Atomic number is either given or must be looked up on the Periodic Table
 Number of protons = Atomic number
 Number of electrons = Number of protons
 Number of neutrons = Atomic mass minus atomic number (n = AM – AN)
Molecular Formulas
chemical formulas: a group of symbols representing the number or type of atoms/ions in a chemical
substance

ex. methane  CH4 (molecular)
ex. salt  NaCl (ionic)
Molecular Formula
 The chemical formula for a binary (two element) molecular compound
 Composed of two non-metals held together by the sharing of electrons forming a covalent bond
NOTE: Non-metals are to the right of the staircase line on the Periodic Table
When Naming Molecular Compounds:
 Prefixes must be used for naming both non-metals (memorize)
 mono: 1
 tri: 3
 penta: 5
 hepta: 7
 di: 2
 tetra: 4
 hexa: 6
 octa: 8
 nona: 9
 deca: 10
 The first name stays the same but the last three letters of the second name change to “ide”
- ex. N2O6  dinitrogen hexaoxide NOT dinitrogen hexaoxygen
 Subscripts indicate prefixes. It also tells the number of atoms in a molecule. If the subscript
for the first name of a compound is (1) then the prefix is omitted.
- ex. CO2  carbon dioxide NOT monocarbon dioxide
* if the name of a compound is given the prefixes become subscripts (e.g) trisulfur tetraoxide
becomes S3O4
 The following groups do not follow these rules: ( Memorize )
 Trivial or Common Names
 methane  CH4 (g)
 ethanol  C2H5OH(ℓ)
 ethane  C2H6 (g)
 water  H2O(ℓ)
 propane  C3H8 (g)
 hydrogen peroxide  H2O2(ℓ)
 butane  C4H10 (g)
 hydrogen sulfide  H2S(g)
 glucose  C6H1206 (s)
 ozone  O3(g)
 sucrose  C12H22011 (s)
 silicon carbide  SiC(s)
 methanol  CH3OH(ℓ)
 ammonia  NH3(g)
 Diatomic Molecules
 All Group VII molecules (H2, F2, Cl2, Br2, I2, At2), oxygen (O2), and nitrogen (N2)
 Polyatomic Molecules
 sulfur  S8 (s)
 phosphorous  P4 (s)
Ionic Compounds
 Formed when a metal is joined to a non-metal (formed when a cation [+
charged atom] is joined to to an anion [- charged atom])
 Held together by an ionic bond caused by the transfer of electrons
Ionic Compounds: ( Properties )
 Are all solids at room temperature
 All dissolve in water to some degree
- ie. Are aqueous (aq)
 Will conduct electricity, but only when dissolved in water
- ie. Are good electrolytes
 Have relatively high melting point and boiling point-higher than
molecular compounds
How to Write Ionic Compounds
 When metals (cations) and non-metals (anions) join, the net
electrical charge of an ionic compound must be zero.
- ie. The sum of all positive charges must equal the sum of all
negative charges
 Never change the charge on an ion from the Periodic Table
- To find the net charge, multiply the charge by using a subscript
 Never Never Never Never use the prefix system
- The first name of the ionic compound stays the same
- The second name of the ionic compound changes the last three letters to “ide”
 All subscripts in an ionic compound must be in lowest terms
NOTE: When the subscript is (1) it is not written in
- When the charges of both the cation and the anion are the same, simply leave out
the subscripts
- When the charges of the cation and the anion are different, write the charge of
the cation as the subscript of the anion and vice versa ( criss-cross )
NOTE: non-metal + non-metal = molecular compound
NOTE: metal(write first) + non-metal(write second) = ionic compound
Multivalent Metals or (the Stock System)
 Multivalent metals are metals that have more than one charge
 Of the charges, the top one listed is the most common. When using
multivalent metals, if no charge is mentioned (if you are not told which
charge to use) use the most common one.
 To write the name of a compound which includes a multivalent metal, you
must write the ion charge (on the metal only) in Roman numerals in
brackets.
 Always start with the anion (non-metal) when given the chem.
formula
 eg. Fe2O3: iron (lll) oxide
eg. SnCl4: tin (lV) chloride
eg. nickel (ll) oxide: NiO
eg. platinum (lV) sulfide: PtS2
NOTE: Know Roman numerals I – X
Polyatomic Ions (Complex Ions)
 A group of atoms that become stable by losing/gaining electrons
 When naming, they are treated identically to monatomic ions
 Put polyatomic ions in brackets when using subscripts
 most end in ( ate ) or ( ite )
 all are negative except for ammonium ( NH4+ )
 eg. lithium sulfate: LiSO4ˉ
eg. calcium nitrate: Ca(NO3)ˉ2
eg. Fe2CO3: iron (l) carbonate
 Complex ions cannot exist alone
 Compounds which contain complex ions are ionic compounds are solids in pure form
but break into ions in water.
NOTES:
Break down of compounds ~
 molecular compound  molecules  atoms
 ionic compound  formula units  ions
Hydrates
 Hydrates are ionic compounds that decompose at relatively low temperatures to produce water
and an associated compound
- eg. CuSO4 • 5H2O(s)
 If heated 5 molecules of water are released and copper (ll)
sulfate remains
 The CuSO4 is now said to be anhydrous (without water)
Naming Hydrates: Two Methods

prefix system
always written as one word
CuSO4 • 5H2O(s) the ‘hydrate’ is the H2O
means ‘in the presence of’
naming
. . . is named as copper (ll) sulfate pentahydrate
 This is the preferred method of
 CuSO4 • 5H2O(s) . . . can also be named as copper (ll) sulfate-5-water
Hydrogen Compounds and Acids
hydrogen compounds:
 Different from most molecular compounds because they form conducting solutions.
Most acids contain hydrogen atoms bonded to some non-metal (eg. HCl) or complex
ion (eg. H2SO4).
Nomenclature of H-Compounds
 Most are named as acids except for the pure compounds (before dissolving)
 The pure compounds are named as though they were ionic
- eg. HCl(g): hydrogen chloride
- eg. HCN(g): hydrogen
cyanide
- eg. H2S(g): hydrogen sulfide
Nomenclature of Acids
 H-compounds must be dissolved in water before an acid can be formed
NOTE: Any solution made when an ionic/molecular compound dissolves in water is
said to be aqueous and the subscript (aq) is added
- eg. HCl(g) + water  HCl(aq)
hydrogen chloride + water  hydrochloric acid
 In naming acids, you first name the H-compound as if it is ionic,
and then convert to the acid name by following these three rules:
IONIC NAME
ACID NAME
 hydrogen (root)ide  hydro(root)ic acid
 hydrogen (root)ate  (root)ic acid
 hydrogen (root)ite  (root)ous acid
NOTE: Two exceptions to these rules are acids which have sulf or phosph as roots.
Syllables are added to these acid names to make them sound better. Thus, H2SO4(aq)
is sulfuric acid not sulfic acid and H3PO4(aq) is phosphoric acid not phosphic acid.
Classification of Acids
binary: if the acid contains a H-compound with hydrogen and one other kind of atom
 eg. HCl(aq): hydrochloric acid
oxo: if the acid contains a H-compound with hydrogen, oxygen, and one other kind of atom
 eg. HClO3(aq): chloric acid
NOTE: If an acid contains a COO- group then the (H) atom is placed at the end of
the group, not in the front
 eg. CH3COOH(aq) (not HCH3COO(aq)): acetic acid (vinegar)
Properties of Acids
1. are solids, liquids, and gases as pure substances at room temperature (like
molecular substances)
2. are soluble in water (like all ionic and some molecular substances)
3. form coloured and colourless solutions (like ionic compounds)
4. form conducting solutions (like ionic compounds)
5. form solutions which turn blue litmus red
6. taste sour and are usually corrosive (burn)
7. pH lower than 7 (0-6.9-acid; 7-neutral;7.1-14-base)
8. will neutralize basic solutions (acid+base = salt +water)
9. react with alkali metals to produce salt and H2 gas
 ex. 2K(s) + 2HCl(aq)  2KCl(s) + 2H2(g)
salt: any ionic compound that is not an acid or a base
Chemical Reactions
 Occur when old bonds within reactants are broken and new bonds within products
are produced
- ie. reactants  products
Why does matter chemically react?
 Atoms want to become stable. They share electrons to form a covalent bond
(molecular compound) and transfer electrons to form an ionic bond (ionic compound).
How can you tell when a chemical reaction occurs?
 You cannot observe it directly, but there are five general observation that
indicate a chemical reaction:
 An unexpected colour change
 A gas is produced
 A solid (precipitate) is produced or dissolved
 A new odour is detected
 A temperature or energy change
- If energy (heat, sound, light, etc.) is absorbed then it is an endothermic
reaction.
- If energy
reaction.
(heat, sound, light, etc.) is released then it is an exothermic
Can a chemical reaction create or use up energy?
 No. The law of conservation of energy states that in any chemical process energy
is neither created nor destroyed, just transformed from one form to another.
- eg. Electrical energy from batteries produces heat and light
Balancing Chemical Equations
Law of Conservation of Atoms: states that in any chemical reaction atoms are
neither created nor destroyed but conserved in the reaction
 The total number of atoms on the reactant side must equal the total number of
atoms on the products side. (Must also have the same type of atoms.)
 3H2O(ℓ)
unit
subscript: gives the number of atoms in a molecule or ions in a formula
coefficient: gives the number of molecules/formula units in a compound
 Always change the coefficient; never change the subscript once it is balanced
for charge
 All coefficients must be in lowest terms and must be whole numbers
Na(s)
+
HOH(ℓ)
NaOH(aq)
+
H2(g)
reactant reacts with reactant to yield/produce product along with product
Na(s) + HOH(ℓ)  NaOH(aq) + H2(g) (unbalanced)
Na – 1
Na - 1
H–2
H-3
O–1
O–1
2Na(s) + 2HOH(ℓ)  2NaOH(aq) + H2(g) (balanced)
Na – 2
Na - 2
H–4
H-4
O–2
O–2

Types of Chemical Reactions
1. Simple Decomposition (SD)
 Occurs when a compound is broken down into it’s elements
 General Form: compound  element1 + element2 + element3 . . .
 eg. hydrogen chloride gas is decomposed into it’s elements
HCl(g)  H2(g) + Cl2(g)
eg. sodium hydroxide is decomposed
2NaOH(s)  2Na(s) + O2(g) + H2(g)
eg. sodium phosphate
4Na3PO4(s)  12Na(s) + P4(s) + 8O2(g)
2. Simple Composition (SC)
 Occurs when elements react with each other to form a compound
 General Form: element1 + element2 + element3 . . .  compound
 eg. glucose is formed from it’s elements
6C(s) + 6H(g) + 3O(g)  C6H12O6(s)
NOTE: Molten means liquid
3. Single Replacement
 Occurs when an element and a compound react to form a different element
and compound
 These reactions involve ions (Remember that a cation must bond with an
anion)
 Write water as HOH(ℓ)
 Use the most common charge when more than one is possible
 General Form: element1 + compound1  element2 + compound2
 eg. 2NaCl(aq) + Br2(ℓ)  2NaBr(aq) + Cl2(g)
eg. 2Na(g) + 2HOH(ℓ)  2NaOH(aq) + H2(g)
4. Double Replacement (DR)
 Occurs when two compounds react and form two new compounds
 The cation of compound1 reacts with the anion of compound2 and the cation
of compound2 reacts with the anion of compound1
 One of the reactants must be (l) or (aq)
 General Form: compound1 + compound2  compound3 + compound4
 eg. 2AgBr(aq) + ZnCl2(l)  2AgCl(g) + ZnBr2(aq)
eg. Pb(NO3)2(aq) + 2KI(s)  PbI2(aq) + 2KNO3(aq)
NOTE: If the state of matter for an ionic reactant is unknown, assume it’s aqueous
NOTE: Coefficients must always be whole numbers
5. Hydrocarbon Combustion
 Occurs when a hydrocarbon (a compound composed of carbon and hydrogen)
reacts with oxygen gas to produce carbon dioxide gas and water vapour
 General Form: hydrocarbon + O2(g)  CO2(g) + H2O(g)
 In general, balance carbon first, hydrogen second, and oxygen third
 eg. methane combusts in the air
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)

eg. 2C2H6 + 7O2(g)  4CO2(g) + 6H2O(g)
6. Other
 Any reaction that is not one of the first five
 Always be given both reactants and products
 eg. CO2(g) + Ca(OH)2(aq)  CaCO3(s) + H2O(ℓ)