science 10 chemistry unit a success review study complete homework
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Science 10 Chemistry Science 10 Chemistry
Unit AUnit A
Significant Digits• Any digit from 1-9 is significant
• Sandwich zeros are significant– E.g. 2.04, 1005.002
• Trailing zeros are significant– E.g. 6.3800, 12 000
• Leading zeros are not significant– E.g. 0.0065
• (+ or -) :– round to the lowest number of decimal
places
• (x or ):– round to the lowest number of sig digs
2.35g + 16.77g + 12.1g = 31.22g
100 g 53.29 g/mol = 1.8765 mol
31.2g
1.88 mol
• counted objects and constants are not
included
WHMIS
•Workplace Hazardous Materials Information System
•WHMIS provides guidelines for handling,
storage and disposal of reactive materials
Know your WHIMIS SYMBOLS!
compressed gas
corrosive flammable and combustible
oxidizing material
poisonous and infectious material causing immediate and serious toxic
effects
poisonous and infectious
causing other toxic effects
Dangerously active material
biohazardous infectious material
Matter
• It is anything that has mass and occupies space
• Has properties which describe the substance
• Two types of properties– physical – chemical
• It can be classified as being a mixture or a pure substance
Substance (Matter)
MixturesPure Substance
Physical Properties
• Describe its appearance or features you can measure:
density color boiling point
melting point
Chemical Properties
• Describes how it reacts with others substances
• It cannot be tested without destroying the substance
E.g. combustion, rusting, decomposition
reactions
Substance (Matter)
Pure Substance Mixtures
Elements Compounds Heterogeneous Homogeneous
Metals
Metalloids
Non-metals
Ionic
MolecularAlloys Solution
s
Classification of Matter
Colloids Suspension
1. Mixtures
• Are created when 2 or more substances combine
• There are 2 types: Heterogeneous Homogenous
A. Heterogeneous Mixture
• 2 or more phases are visible, or two or more substances are visible
• E.g. chicken soup, pulpy orange juice
• There are two types: Suspension Colloid
Suspension
• Is a mixture in which the components are in different states
• E.g. Mud (water & sand)
Colloids
• A mixture in which the components can’t be easily separated
• E.g. Milk or salad dressing
B. Homogeneous Mixture
• A mixture of 2 or more substances that appear as one….( uniform properties)
• E.g. kool-aid, coffee, air
• There are 2 types: Alloys Solutions
Alloys
• A homogenous mixture of 2 or more metals
• E.g. brass – copper-zinc
Solutions
• A mixture where you can not see the different parts
• The composition does not change
• They can be: Elements Compounds
2. Pure Substances
A. Elements
• Cannot be broken down further
• contains only one type of atom
• E.g. Gold
• Have 3 categories: Metals Nonmetals metalloids
Metals• Are:
ductile (can be stretched into wire)
lustrous (shiny)
conductors
malleable
found to the left of the staircase line
make up approx. 80% of the elements
silver
Non-Metals
• Are:
non-ductile
dull
non-conductors
brittle
found to the right of the staircase line
make up approx. 20% of the elements
Sulphur
Metalloids
• Are elements that have properties of both metals and non-metals
E.g. carbon – dull, conducts
silicon – lustrous, non- conductor
arsenic
silicon
B. Compounds
• Are two or more elements bonded together in fixed proportions
• E.g. H2O , C6H12O6
The Periodic Table
• It was developed by Dmitri Mendeleev in the mid 1800’s
Atomic Atomic MassMass
Element Element NameName
Element Element SymbolSymbol
Atomic Atomic NumberNumber
Ionic Ionic ChargeCharge
• It follows certain rules (patterns) called the “Periodic Law”
• It is arranged into groups (families) and periods
Period
• Are the horizontal rows on the periodic table
Period 1 - 2 elements
Period 2 - 8 elements
Period 3 - 8 elements
Period 4 - 18 elements
Period 5 - 18 elements
Period 6 – 32 elements
Period 7 – 32 elements
Groups (Families)
• Vertical columns
• Have similar properties
(behave in a similar manner)
• The number indicates the # of electrons in the last energy level
• Have 2 numbering systems :
a) Arabic #’s (1 -> 18)
b)Roman numerals & letters (IIA ,VIIB)
Group IA
• Alkali Metals (E.g. Li, Na, K)
• most reactive metals
• never found in pure form in nature
• all have one electron in the last energy level
Group IIA
• Alkaline Earths (E.g. Mg, Ca, Ba)
• very reactive metals
• all have 2 electrons in the last energy level
Group VII A
• Halogens E.g. (Cl2, Br2, )
• most reactive non-metals
• all are diatomic elements E.g.( F2, I2, At2 )
• are one electron short of a full energy level
Group VIII A
• Are Noble Gases
• non-reactive elements (inert gases)
• all have full energy levels
Series• Are at the bottom of the periodic table
• Lanthanum (Lathanides) are rare earth metals located 57-71
• Actinium (Actinides) Are located 89 to 103
Transition Elements
• Belong to Group B
• Are located in the middle of the table
magnetic elements are found here
Trends of Transition Metals
• metals:– get more reactive as you move
(R L) and
– The most reactive metal is FRANCIUM
• Non-metals :
– get more reactive as you move (L R) and
– this excludes noble gases
Atomic Structure
• Describes what an atom is made up of
• http://www.youtube.com/watch?v=07yDiELe83Y&feature=related
• A Theory was proposed by Dalton in 1808
Dalton’s Theory1. All matter is composed of tiny,
indivisible particles called atoms.
2. Atoms of an element have identical properties
3. Atoms of different elements have different properties
4. Atoms of two or more elements can combine in constant ratio to form new substances.
Know this: Page 22 Textbook
J.J Thompson: 1897• credited with discovery of
electrons
• “raisin bun” model or “plum pudding” model
• atom is a sphere which is positive, with negative electrons embedded in it like raisins in a bun
• most of the mass is associated with the positive charge
Ernest Rutherford: 1911
• atoms have a nucleus which is positive and has most of the mass
• most of the atom is empty space occupied by the moving negatively charged electrons
• proposed the existence of protons
http://www.learnerstv.com/animation/animation.php?ani=121&cat=chemistry
Neils Bohr: 1913 • electrons move in circular orbits around the
nucleus
• cannot exist between orbits
James Chadwick: 1932• showed that the nucleus must contain heavy
neutral particles to account for all of the atom’s mass (neutrons)
Schrodinger/de Broglie: 1930
• quantum mechanical model
• electrons have distinct energy levels
• exact locations of electrons are not defined, but the probable location in a region of space can be predicted
Atoms• Are the smallest part of an element
which retains the chemical and physical properties of an element.
• They are neutral
• It is made up of 3 sub-atomic particles:
–Protons
–Neutrons
–Electrons
Protons (P+)• Are large positively charged particle in the
atom’s nucleus
• They make up 99% of the mass of the atom
• the number of protons determines the atomic number
• E.g. Cu has 29 protons… atomic number =29
Neutrons (N0)
• Are large particles in the atom’s nucleus
• They hold the nucleus together
• They have no charge
Electrons (e-)• Are the smallest particle in an
atom
• They have a negative charge
• Are located outside the nucleus of the atom
• They take up most of the space
• They are arranged in energy levels
• The # of electrons in each level is :
• Level # of e-
• 1 2
• 2 8
• 3 8
• 4 18
You MUST MEMORIZE
THIS!!!!!
Atomic Number
• Is the # of protons in an atomE.g. oxygen:
atomic # = 8 it has 8 protons
Atomic Atomic MassMass
Element Element NameName
Element Element SymbolSymbol
Atomic Atomic NumberNumber
Ionic Ionic ChargeCharge
Mass Number
• Also called the atomic mass
• its the sum of the protons and neutrons (These are averages)
Atomic Atomic MassMass
Element Element NameName
Element Element SymbolSymbol
Atomic Atomic NumberNumber
Ionic Ionic ChargeCharge
• E.g. Lithium Atomic Number (so it has 3 Neutrons) Atomic Mass (round it)
• Try: Mercury Atomic Number ___ (so it has _ Neutrons) Atomic Mass ___ (round it)
• Try: Chlorine Atomic Number ___ (so it has _ Neutrons) Atomic Mass ___ (round it)
3
6.9
Finding The Number of Neutrons
• Use the Formula: Neutrons = Atomic Mass – Protons
• E.g. Oxygen Mass = ____ Protons = ____
Neutrons = Atomic Mass – Protons Neutrons = _____ - ____ Neutrons = _____
Try the Following
• E.g. Calcium Mass = ____ Protons = ____
• Neutrons = Atomic Mass – Protons
• Neutrons = _____ - ____
• Neutrons = _____
Isotopes
• Have a different atomic mass
• They have a different number of neutrons…. but same number of protons
• E.g. Copper has 35 neutrons ..but an isotope of copper can have 36
Isotope Notation
x Az
A = symbolx = mass # (p+ + n°)z = atomic # (p+)
Try copper – 64 (together)
Isotope Notation x
Az
A = symbolx = mass # (#p+ + n°)z = atomic # (#p+)
copper - 64
__ Cu__
# p+ = ____#n° = ___ - ___ =___# e- = ____
Try copper - 62
# p+ = _____#n° = _____# e- = _____
Try the Following
• Write the isotope notation for the following:
Cobalt - 61 Cobalt – 60
EELR - Electron Energy Level Representations
• It is a diagram that shows the following:
– The Nucleus
– Energy Levels
– Valence electrons
• Energy Levels: show the number of electrons in each
level
number of levels = the period #
• Valence Electrons: Are electrons in the last energy level
It is the group #
E.g. Sodium - atomic # ___ - mass # ___
# p+ = # e- =#n° =
____________ - ____ = _____ (round) ____
__ e-
__ e-
__ e-
____ e-
Level 1 = 2Level 2 = 8 Level 3 = 8
Level 4 = 18
E.g. Sodium - atomic # 11 - mass # 22.99
# p+ = # e- =#n° =
111122.99 - 11 = 11.99 12
2 e-
8 e-
1 e-
11 e-
Level 1 = 2Level 2 = 8 Level 3 = 8
Level 4 = 18
Krypton - atomic # _____- mass # _____
# p+ = # e- =#n° =
_________ - ____ = ____ round ___
___ e-
___ e-
___ e-
___ e-
____ e-
Level 1 = 2Level 2 = 8 Level 3 = 8
Level 4 = 18
Krypton - atomic # 36- mass # 83.80
# p+ = # e- =#n° =
363683.80 - 36 = 47.80 48
2 e-
8 e-
8 e-
18 e-
36 e-
Level 1 = 2Level 2 = 8 Level 3 = 8
Level 4 = 18
Assignment
Draw an EELR for each of the following atoms
(a) berylium(c) chlorine(e) calcium(f) magnesium
Ions
• Are atoms or group of atoms that have:– become electrically charged – gained or lost electrons to be isoelectronic with
a noble gas (same configuration)
• They can be:
– Monatomic
– Polyatomic
Their outer orbital is filled
A. Monatomic Ions
• Are made of one element
• Can be:
– Cations
– Anions
Cations
• Are positively charged
• Are METALS
• lose electrons …….to obtain a stable electron configuration (like a noble gas)
• E.g. Sodium ( Na+1 )
• E.g. Sodium ion• Losses 1 electron• It tries to have the same configuration as Neon• Na+1 has 10 electrons
• Try: Calcium Ion• Losses _____ electron(s)• It tries to have the same configuration as _____• Ca 2+ has ________ electrons
Anions
• Are negatively charged
• They are NON-METALS
• gain electrons….. to obtain a stable electron configuration (like a noble gas)
• They have an “ide” ending
• E.g. Sulfide (S 2-)
• E.g. Fluorine ion• Gains 1 electron• It tries to have the same configuration as Neon• F-1 has 10 electrons
• Try: Oxygen Ion• Gains_____ electron(s)• It tries to have the same configuration as _____• O2- has ________ electrons
B. Polyatomic Ions
• Are a group of atoms that act like anions. (Except: ammonium…it acts like a cation)
• They are located at the top of the periodic table
They don’t use bi anymore… use hydrogen instead
Different Ways of Writing Acetate
CH3COO− or C2H3O−
EELR of Ions
• It is a diagram that shows the following:
– The Nucleus
– Energy Levels
– Valence electrons
• The only difference is the electrons!!!
Electrons = # of protons - the charge
Nitride - atomic # _____- mass # _____
# p+ = # e- =#n° =
________ - (-___) = ________ - ___ = ____ ______
___ e-
___ e- ____ e-
N3-
Nitride - atomic # 7- mass # 14.01
# p+ = # e- =#n° =
77 - (-3) = 1014.01 - 7 = 7.01 7
2 e-
8 e- 10 e-
N3-
Cadmium ion - atomic # ____- mass # ____
# p+ = # e- =#n° =
________ - (+___) = _______ - ____ = _____ ___
___ e-___e-
___ e-___e-___e-
___ e-
Cd 2+
Cadmium ion - atomic # 48- mass # 112.41
# p+ = # e- =#n° =
4848 - (+2) = 46112.41 - 48 = 64.41 64
2 e-8 e-
46 e-8 e-18 e-
10 e-
Cd 2+
Iron (III) ion - atomic # ___- mass # ___
# p+ = # e- =#n° =
________ - (+___) = _______ - ___ = _____ ==> ____
___ e-__ e-
____ e-__ e-__e-
Fe3+
Iron (III) ion - atomic # 26- mass # 55.85
# p+ = # e- =#n° =
2626 - (+3) = 2355.85 - 26 = 29.85 ==> 30
2 e-8 e-
23 e-8 e- 5e-
Fe3+
Draw an EELR for each of the following ions:
1. Sulfide2. strontium ion3. Copper (II) ion4. Iodide5. Vanadium (V) ion
Elements
• Metallic elements:– exist as single atoms (monatomic)
– The formula is the symbol followed by the state
E.g. [sodium] Na (s)
• Non-metals:
– Excluding noble gases don’t exist as single atoms… they are diatomic or polyatomic
H2
N2 O2F2
P4S8 Cl2
Br2
I2polyatomic
Chemical Bonds
• Are interactions that occur between atoms or molecules
• There are two types: Covalent Ionic
Covalent
• Are formed when electrons are shared between atoms
• It occurs between non-metals
1+ 1+
Each needs one electron to complete it’s 1st orbital
Hydrogen Ions
1+ 1+
Single Bond
Hydrogen Molecule
H -- H
H2
Ionic
• Are formed when electrons TRANSFER from one atom to another
• It occurs between a metal and a nonmetal.
Compounds
• are two or more substances that are held together by chemical bonds
• There are two types:
– Molecular
– Ionic
Molecular
• Are solid, liquid or gas at room temperature
• Contain only non- metals
• Have a covalent bond
• Don’t conduct electricity
• may dissolve in water to produce either
(a) neutral molecular solution
(b) acidic solution
Ionic Compounds
• Are crystal solids at room temperature
• occur because of a force of a attraction between the + and - ions
• contain a metal and a non metal
• have an ionic bond
• Have high melting point and boiling points
• Conduct electricity in water
• May dissolve in water to produce either a
(a)neutral ionic solution
(b)basic solution.
Ionic & Molecular CompoundsIonic Molecular
cation + anion (metal + non-metal)
all elements are non -metals
NO PREFIXES PREFIXES
ALL SOLIDS AT R.T. SOLIDS, LIQUIDS & GAS
Solutions conduct electricity
solutions do not conduct electricity
solutions are basic or neutral
solutions are acidic or neutral
ex. NaCl ex. OF2
Identify if it is M or I
• A yellow gas forms a neutral solution_____
• A purple solid dissolves in water to produce a conducting solution _____
• A white solid dissolves in water and turns red litmus paper blue ______
(solutions)
• A yellow gas forms a neutral solution____M_
• A purple solid dissolves in water to produce a conducting solution __I___
• A white solid dissolves in water and turns red litmus paper blue __I____
Naming Naming &&
Writing FormulasWriting FormulasForFor
Molecular & Ionic CompoundsMolecular & Ionic Compounds
Naming Molecular CompoundsNon-metal + Non-metal
• Step 1: write the element name for the first
non-metal add a prefix ….(if there is more than
one atom)
• Step 2: Write the second element name and
add an ide ending Add a prefix
Exceptions
• DO NOT use any prefixes at all if the first element is hydrogen ……these are acids
Prefixes
1 = 6 =
2 = 7 =
3 = 8 =
4 = 9 =
5 = 10 =
mono
ditri
tetra
penta
hexa
hepta
octa
nona
deca
Examples
• P4O10 tetraphosphorus decaoxide
• bromine heptahydride
• H2S hydrogen sulfide
BrH7
Try the Following
• CO(g)
• CO2(g)
• N3F8
carbon monoxide
carbon dioxide
Trinitrogen octafluoride
Writing Formulas for Molecular Compounds
• Steps: Write each elements symbol Write the subscript number (the prefix)
• E.g. dinitrogen oxide N2O
disulfur oxide S20
Try the Following
• oxygen dibromide
• diphosphorus pentasulphide
• carbon tetraiodide
• phosphorus pentachloride
OBr2
P2S5
CI4
PCl5
Molecular Compounds that Must be memorized !!!
ammonia
water
NH3 ( g) =
H2O ( l) =H2S ( g) =
CH4 ( g) =
CH3OH ( l) =C2H6 ( g) =C2H5OH ( l) =
C6H12O6 ( s) =
hydrogen sulphide
methane
methanolethane
ethanol
glucosesucroseC12H22O11 ( s) =
hydrogen peroxideozoneO3 ( g) =
H2O2 ( l) =
Naming Binary Ionic CompoundsMetal + non-metal
• DO NOT USE PREFIXES
• Steps
1.Write the metal 1st
2. Write the non-metal 2nd with an ide ending
E.g. NaF
Na2Ssodium fluoride
sodium sulphide
two sodium ions are bonded with one sulphide ion… this doesn’t matter for naming ionic compound
Try the Following
• LiF
• KCl
• BeS
• Rb3P
• MgF2
• Na2O
• CsBr
lithium fluoride
potassium chloride
beryllium sulphide
rubidium phosphidemagnesium fluoride
sodium oxide
cesium bromide
Try the Following
• KCl
• MgBr2
• Ba3N2
• ScP
potassium chloride
magnesium bromide
barium nitride
scandium phosphide
Writing Formulas for Binary Ionic Compounds
• Steps
1.Look up the symbol for each…& write the metal first
2.Balance the charges (total + charges = total – charges)
3.Use subscripts to show the # of each element
sodium oxide
1+ Charge 2 Charge
1+ 2 = 2+ 2 1 = 2
Na2O
calcium phosphide
2+ Charge 3 Charge
2+ 3 = 6+ 3 2 = 6
Ca3P2
Try the Following
• magnesium chloride
• calcium chloride
• zinc sulphide
• silver sulphide
• germanium oxide
• calcium arsenide
• magnesium nitride
MgCl2
CaCl2
ZnS
Ag2S
GeO2
Ca3As2
Mg3N2
Try the Following
• lithium iodide
• zinc fluoride
• strontium phosphide
• silver oxide
• germanium arsenide
LiI
ZnF2
Sr3P2
Ag2O
Ge3As4
Naming Multivalent Ionic Compounds
• Transition metal ions have more than one possible charge
Cu2+, Cu+, Fe3+, Fe2+
transition metal + non-metal
• Steps
1. Write metal 1st with the charge in roman numerals
2.Write non-metal second
remember the charges have to balance
Roman Numerals (I,II,III,IV,V,VI,VII)
Examples
uranium (VI) fluoride
chromium (III) nitride
cobalt (II) chloride
U6+ F–
Cr3+ N3-
Co2+ Cl-
UF6
CrN
CoCl2
Try the Following
• AuBr
• CrCl2
• Co2O3
• VS2
• PuN2
gold (I) bromide
chromium (II) chloride
cobalt (III) oxide
vanadium (IV) sulphide
plutonium (VI) nitride
Naming Complex Ions
Metal + complex ion
• Steps:
1. Name the metal ion
2. Name the complex ion E.g.) PO43
Note: NH4+ (ammonium ion) is the
only positive complex ion…it will take the place of a metal
Examples
• CaCO3
• Ba(OH)2
• (NH4)3N
• CaCO3
• Ba(OH)2
• (NH4)3N
Solutions
Calcium carbonate
Barium hydroxide
Ammonium nitride
Try the Following
• KIO3
• NaCH3COO
• MgSO3
• NH4NO3
• Ca3(PO4)2
potassium iodate
sodium acetate
magnesium sulphite
ammonium nitrate
calcium phosphate
Writing Formulas For Complex Ions
• Steps:
1. Look up the symbol for each ion
2. Balance the charges
• Note: if you need more than 1 complex ion to balance the charges use brackets
• E.g. Ca(CH3COO)2
2+ 1 -
Try the Following
• aluminum phosphate
• calcium sulphite
• scandium acetate
• ammonium sulphate
• nickel (II) phosphate
• aluminum chlorate
AlPO4
Al(ClO3)3
CaSO3
Sc(CH3COO)3
(NH4)2SO4
Ni3(PO4) 2
SolubilitySolubility
Will the compound dissolve in water?Will the compound dissolve in water?
Soluble • Refers to whether or not the compound
dissolves in water
• If it is…. the compound is aqueous (aq)
• All acids are soluble
• Some ionic compounds are soluble… the rest are solids
Is It soluble?• This will apply to ionic compounds (only)
• Steps
1. Find each ion in the boxes across the top
2. if it is soluble it will have (aq) aqueous
3. If it does not dissolve it will have (s) solid.
Determine if the following compounds are soluble in water. Use the proper subscript to indicate the state.
• AgCl• BaCO3
• LiOH• Ca2(PO4)3
• NaCl• CaI2
• Pb(NO3)2
• HMnO4
• AgCl
• BaCO3
• LiOH
• Ca2(PO4)3
• NaCl
• CaI2
• Pb(NO3)2
• HMnO4
(s)
(s)
(aq)
(s)
(aq)
(aq)
(aq)
(aq)
Acids & BasesAcids & Bases
Acids
• They are always soluble in water
• Conduct electricity
• Taste sour
• React with metals to produce hydrogen gas (H2(g))
• Neutralize a base
• they ALWAYS have hydrogen ….usually as the first element
• E.g. HCl(aq) , H3PO4(aq)
• There are three types: Binary Oxo Organic
Binary Acids
• contain only H and one other element (Cl, Br, etc.)
• E.g. HCl (aq)
Oxo Acids
• Contain H and Oxygen
• E.g. H3PO4(aq)
Organic Acids
• Contain C, H, & O
• The H is written at the end
• All have COO-
• E.g. CH3COOH(aq) – acetic acid
C6H5COOH(aq) - benzoic acid
HOOCCOOH(aq) – oxalic acid
Acid Indicators
• Turns blue litmus paper red
• Able to turn bromothymol blue to yellow
• Phenolphthalein remains colorless
• E.g. lemon juice
Bases
• Are usually soluble in water
• Conduct electricity (not weak ones)
• Neutralize acids
• Taste bitter
• Usually solids
• Feel slippery
Base Indicators
• Turns red litmus paper blue
• Bromothymol blue remains blue
• Turns phenolphthalein pink
• E.g. baking soda, Rolaids, soap, Draino crystals
Naming Acids
• Steps:
Hydrogen ____ide becomes hydro____ic acid
Hydrogen ____ate becomes _______ic acid
Hydrogen ____ ite becomes _______ ous acid
Examples
• HF (aq) • H2SO3 (aq)
• H3BO3 (aq)
• HCl (g)
hydrogen fluoride = hydrofluoric acid
hydrogen sulphite = sulphurous acid
hydrogen borate = boric acid
Hydrogen chloride (not an acid)
hydrosulphuric acid phosphorus acid
carbonic acid
hydrogen phosphite H3PO3 (aq)
1+ 3-
hydrogen carbonate H2CO3(aq)
1+ 2-
hydrogen sulphide H2S (aq)
1+ 2-
Try the Following
Writing Acid Formulas• Steps:
1. Use the naming rules in the opposite direction
• Example:
hydrosulphuric acid
hydrogen sulphide H2S(aq)
Try the Following
• carbonic acid
• chlorous acid
hydrogen carbonate
hydrogen chlorite
H2CO3(aq)
HClO2(aq)
Naming Bases
• Steps:
1. Write the metal name 1st
2. Write hydroxide or bicarbonate
E. g. NaOH sodium hydroxide
Try the Following
• KOH
• Ba(OH)2
• NaHCO3
potassium hydroxide
barium hydroxidesodium bicarbonate
Chemical Reactions• Can cause a physical or a chemical
change
• Always results in the formation of a new substance
• Evidence:
1. Temperature change
2. Formation of a precipitate
3. Colour change
4. Gas produces
states states
balancing
Reactants Products
1 H2 (g)+1 ZnCl2(aq)2 HCl(aq)1 Zn(s)+
Energy Changes
• Can occur in the form of heat, light, electrical, or mechanical
• There are two types: Endothermic Exothermic
• Endothermic Energy is absorbed (enters) Reactants + Energy products
• Exothermic Energy is released (leaves) Reactants product + energy
Balancing Equations
• There must be equal numbers of each element on both sides of the equation
Use lowest numbers
Example
____Mg(s) + ___ O2(g) ___ MgO(s)
____ H2O(l) ____H2(g) + ____ O2(g)
• When chemicals react they follow the Law of Conservation of Matter:
Matter can not be created or destroyed it only changes form
• Mass of reactants = mass of products
Counting Practice
• How many of each element are in the following compounds?
1. NaCl 5. NH4CH3COO
2. BaBr2 6. 3 (NH4)2S
3. (NH4)3P 7. 2 CaCl2
4. Ba(OH)2 8. 8 PbI2
9. 4 Zn(CH3COO)2
Balancing Practice
____ Cu(s) + ____ AgNO3(aq) ____ Ag(s) + ___ Cu(NO3)2(g)
___ Cl2(g) + ____ NaBr(aq) ____ Br2(l) + ____ NaCl(aq)
____ KI(aq) + ____ Pb(NO3)2(aq) ___ PbI2(s) + ___NO3(aq)
____ CH4(g) + ____ O2(g) ____ CO2(g) + ____ H2O(g)
Types of Reactions
• There are 5 types of reactions: Simple Composition Simple decomposition Single replacement Double replacement Hydrocarbon combustion
Composition/ Formation Reactions
• Elements combine to form a compound
Element + element compound (s)
• These are usually exothermic
E.g. 2 Mg (s) + O2 (g) 2 MgO(s)
Try the FollowingMg(s) + Cl2(g)
2+ 1-
Fe(s) + O2(g) 3+ 2-
MgCl2(s)
Fe2O3(s)4 3 2
Zn(s) + O2 (g) 2+ 2-
ZnO(s)2 2
Simple Decomposition
• Compound decomposes into its elements
Compound element + element
• These are usually Endothermic
E.g. 2 H2O(l) 2 H2(g) + O2(g)
Try the Following
HMnO4(s) H2(g) + Mn(s) + O2(g)2 2
HCl(g)
4
H2(g) + Cl2(g)2
Pb(NO3)2(s) Pb(s) + N2(g) + O2(g) 3
Single Replacement
• An element reacts with an ionic compound to form a different element and compound
• This occurs in water (use solubility chart)
element + compound element + compound
E.g. Cu(s) + 2 AgNO3(aq) 2 Ag(s) + Cu(NO3)2(aq)
Example
Cl2(g) + 2 NaBr(aq) Br2(l) + 2 NaCl(aq)
Zn (s) + NaCl (aq) ZnCl2 + Na2 2(aq)
Pb (s) + Cu(NO3)2 (aq) Pb(NO3)2 + Cu(aq)
Mg + HOH Mg(OH)2+ H22 (s)
KCl + Br2 KBr + Cl2(l)
(l)
(aq) (aq)2 2
Try the Following
(g)
(g)
(s)
(s)
(s)
Double Replacement
• Two ionic compounds react to form two different ionic compounds
• It occurs in water (solubility chart)
• There are two types of reactions: Precipitation Neutralization
compound + compound compound + compound
a. Precipitation Reaction
• One product that is formed is insoluble
E.g. Pb(NO3)2(aq) + 2 KI(aq) PbI2(s) + 2 KNO3(aq)
compound + compound insoluble + compound compound
solid
b. Neutralization
• An acid reacts with a base to form water and salt
acid + base water + salt
E.g. 1 HCl(aq) + 1 NaOH(aq) 1 H2O(l) + 1 NaCl(aq)
Hydrocarbon Combustion• A hydrocarbon is made up of C &H (E.g. CH4)
• Occurs when a hydrocarbon burns in the presence of oxygen
• it always produces CO2 (g) + H2O (g)
C?H? + O2(g) CO2(g) + H2O(g)
E. g. CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)
hydrocarbon
C6H12O6(s) + O2(g) CO2(g) + H2O(g)6 66
CO2(g) + H2O(g)C2H6 (g) + O2(g)
C2H6 (g) + O2(g) CO2(g) + H2O(g)
2 3
2 7 4 6
CO2(g) + H2O(g)C10H22 (l) + O2(g) 2 31 20 22
Try the Following
Other • any reaction that does not follow any of the above patterns
E.g. Ca(OH)2(aq) + CO2(g) CaCO3(s) + H2O(g)
HBrO3(aq) + HBr HOH(l) + Br2(l)
35 3
Mole
• is a quantity equal to 6.02 x1023 atoms, ions, molecules, etc…. Which is Avogadro’s number
• originally defined as the number of atoms in exactly 12 g of carbon-12
• it is used to determine the mass of all other elements
Molar Mass • Is the mass of 1 mole of a substance
Symbol – M Units – g/mol
Steps: 1. Write the correct chemical formula. 2. List all elements present3. Determine how many of each element are
present4. Multiply by the atomic molar mass of that
element. 5. Find the sum.
Example 1:
Determine the molar mass of barium hydroxide
Ba(OH)2
BaOH
122
x 137.33 g/mol x 16.00 g/molx 1.01 g/mol
= 137.33 g/mol= 32.00 g/mol = 2.02 g/mol 171.35 g/mol
Example 2:
Determine the molar mass of magnesium nitrate
Mg(NO3)2
MgNO
126
x 24.31 g/mol x 14.01 g/molx 16.00 g/mol
= 24.31 g/mol= 28.02 g/mol = 96.00 g/mol 148.33 g/mol
Finding the # of Moles
n = m M
n = moles (mol)
m = mass (g)
M = molar mass (g/mol)
m = nM
E.g.1 How many moles are there in a 25.0 g sample of lithium sulfate?
Li2SO4 = Li – 2 x 6.94 g/mol
S - 1 x 32.06 g/mol
O – 4 x 16.00 g/mol
109.94 g/mol
n = m/M
= 25.0 g 109.94 g/mol = 0.227 mol
E.g.2 How many moles are there in a 32.8 g sample of potassium permanganate?
KMnO4 = K – 1 x 39.10 g/mol Mn - 1 x 54.94 g/mol O – 4 x 16.00 g/mol
159.04 g/mol
n = m/M
= 32.8 g 159.04 g/mol
= 0.20754… mol = 0.208 mol
Example 3
What mass of copper (II) sulfate is present in a 0.3750 mol sample?
m = ? n = 0.3750 mol M =
CuSO4 Cu – 1 x 63.55 g/mol S - 1 x 32.06 g/mol O - 4 x 16.00 g/mol 159.61 g/mol
m = nM = (0.3750 mol)(159.61 g/mol) = 59.85 g
Example 4
What mass of sodium hydrogen carbonate is present in a 2.50 mol sample?
m = ? n = 2.50 mol M =
NaHCO3
Na - 1 x 22.99 g/molH - 1 x 1.01 g/mol C - 1 x 12.01 g/mol O - 3 x 16.00 g/mol 84.01 g/mol
m = nM = (2.50 mol)(84.01 g/mol) = 210g