science 10 chemistry unit a success review study complete homework

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