Vraj Rao - Mr. Z. Syed - SCH4U Grade 12 Chemistry

Unit 1:
Structures and Properties of Matter

History of the Atom

Empedocles

Proposed that Matter was made of 4 elements (fire, air, earth, water)

Democretes

Thought experiment

"What would happen if matter could be cut in half an infinite number of times?"

After each cut, the identity would be unchanged

Eventually reach a piece that cannot be cut

Atomos

Lavosier

Law of Conservation of Mass

Mass is neither created or destroyed, only conserved

Proost

Law of Definite Proportions

Different samples of any pure compound contain the same elements in the same proportions by mass

Dalton

Law of Multiple Proportions

The masses of one element that can combine chemically with a fixed mass of another element are in a ratio of small wholes

Atomic Theory

Matter is composed of indivisible particles

All atoms of a particular element are identical

Different elements have different atoms

Subatomic Particles

Electrons

1897 by J.J. Thomson

While working with cathode rays

Observed how they rays were bent by magnets, and how it could help estimate the mass of the rays

Concluded that the rays were in fact small negatively charged particles - called them Corpuscles (we now call them electrons)

Went on the propose the Plum Pudding Model

Protons

Rutherford adjusted his experiment

Bombarded "N" atoms, and noticed that Hydrogen ions were released

Concluded that Hydrogen ions are fundamental particles - called them Protons

Light

A type of electromagnetic radiation (ER)

Carries energy through space

Many types of ER in addition to visible light

Electromagnetic Spectrum

Shows the types of ER arranged in order of decreasing wavelength

Wavelength and Frequency

Pattern of peaks and troughs

Peak: Top of pattern

Trough: Botom of pattern

Frequency: The number of cycles (WL) per second that passes a given point

Terminology

Refraction

Whenever light goes from 1 medum to another, at an angle, the angle changes, making the light beam beng

Dispersion

The amount of bending that occurs depends on WL of light

As WL increases, bending decreases

Planck's Constant

Matter absorbed/emitted energy in whole number multiples of hv

Shows that energy can only be transferred in "packets" AKA quantum (plural is quanta)

Photoelectric Effect

1887 by Heinrich Hertz

Certain metallic surfaces lose their negative charges when exposed to light/electricity (not the light's intensity)

Frequency of light determines how quickly the metal would lose its charge

Einstein later explained

Light itself consists of individual quantum particles, called Photons

Compton Effect

Loss of frequency from shattered X-ray suggests that X-rays have momentum, because light can have properties and behave like Matter

Emission Spectrum

1880s by J.J. Balmer

Specific WL given off by the light of atoms

WL = B (n^2/n^2-m^2)

WL = wavelength of the light emitted
B = Balmer constant for Hydrogen = 364.50682 m
n = integer such that n>m
m = 2

Atoms

Bohr Model

Electrons orbit the nucleus at certain discrete distances

Orbits have defined energies called shells/levels

Energy of shells increase as one proceeds away from the nucleus. Electrons only gain/lose energy by jumping from an allowed orbit to another

Specific quantities of energy/light are required for an electron to "jump" to a higher orbit - AKA - specific elements absorb specific quantities for energy/light

Emission Spectrum = Absorption Spectrum

Energy released when an electron falls from level 4 to 1 is the same amount of energy required for an electron to jump from level 1 to 4

delta Ejump = delta Efall

Binding Energy

Rydberg's Constant (Rh) = 2,18^10^-18 J

E = Rh/n^2

Structures

Lewis Dot Diagrams

Write the chemical name of the substance and draw its valence electrons around it in a circle. Keep in mind the sharing of electron pairs in compounds

VSEPR

Steps

#1. Draw Lewis-Dot-Diagram for the molecule and count the number of electron charge clouds surrounding the atom

#2. Identify the charge clouds as bonding electrons or as lone pairs

#3. Use VSEPR Notation to determine the molecule's shape

AXmEn

A = Central atom
X = Bonding set
m = Number of bonding sets
E = Lone pairs
n = Number of lone pairs

Molecular Shape and Polarity

Bonds

Types

Ionic

Polar Covalent

Non-Polar Covalent

A molecule may have polar bonds but may not be polar

Use electronegativity to predict the polarity of each bond

Steps

#1. Draw the Lewis structure

#2. Draw VSEPR diagram based on Lewis structure

#3. Add electronegativity of the atoms and assign (lamda+) and (lamda-) to the bonds

#4. Draw in teh bond dipoles

Polarity

When 2 atoms bond

Sharing of a pair of electrons can be polar, non-polar, or ionic

When bonding is polar

a bond dipole is created

The more EN atom has a partial negative charge

The less EN atom has a partial positive charge

Forces

Intramolecular

Forces excerted within a molecule/polyatomic ion

Intermolecular

Forces exerted between molecules/polyatomic ions

Influence physical properties of substances

Forces of attraction/repulsion between molecules

4 Categories

Dipole-Dipole

Attraction between opposite partial charges of polar molecules

Polar molecules are more attracted to each other than similar N-P molecules

Contribute to higher melting/boiling points in polar molecules

Ion-Dipole

Attraction between partial charges of polar molecules/ions

Can occur between

Polar molecule (negative end) and cation

Cations usually smaller than anions

Polar molecule (positive end) and anion

Dipole-Induced Dipole

Electrons in atoms are in constant motion

Possible to induce formation of dipoles in NP molecules

Attraction between a polar molecule and a temp. dipole of a NP molecule

Attraction between an ion and a temp. dipole of a NP molecule

Dispersion

Weak attraction between all molecules

Occurs because NP molecules spontaneously form temp. dipoles

Strongest to Weakest

#1. Hydrogen bonding

#2. Dipole-Dipole interactions

#3. Ion-Dipole interactions

#4. Dipole-Induced Dipole interactions

#5. Dispersion forces

Hybridization

Valence Bond Theory

Covalent bond formation and molecular shapes based on orbital overlap

Molecular Orbital Theory

Covalent bond formation and molecular shapes based on the formation of new molecular orbitals

Principals

#1. Region of orbital overlap has a MAX of 2 electrons

#2. Should be MAX orbital overlap

#3. Concept of atomic orbital hybridization is used

Types

Single Bond

Involves 1 SIGMA bond

Double Bond

Involves 1 SIGMA bond and other PI bonds

Triple Bond

Involves 1 SIGMA and 2 PI bonds

sp = Linear
sp^2 = Trig. Planar
sp^3 = Tetrahedral
sp^3d = Trig. Bypyram.
sp^3d^2 = Octahedral
ETC

Number of hybrid orbitals that form = Number of atomic orbitals that combined to make the hybrid orbitals

Unit 2:
Organic Chemistry

Organic Molecules

Naming Hydrocarbons

Components of Names:

Root: The number of Carbons in
the largest chain

Root Names:

*insert pic*

Suffix: The functional group

Suffix Names:

*insert pic*

Basic Naming Rules

#1. Identify the suffix

#2. Identify the root (longest continuous
chain of Carbons)

#3. Assign numbers (locants) to the principal functional
group as lowest common chain

Functional Groups

Types

Alkanes (-ANE)

Hydrocarbons with single bonds

Steps

#1. Find the longest C chain and number them appropriately in ascending order (find the correct name from the chart)

#2. Find all the substituents and name them correctly from the chart

#3. Identify which numbers of substituents are connected to the C chain

Other "Add-ons"

Cycloalkane

Ring like structure of alkanes

Haloalkanes

Alkanes with halogen atoms

Structural Isomers

Compounds with the same molecular formula but with different bonding arrangements

Aromatic Hydrocarbons

Contain a benzene ring as a base

ORTHO

O-

Unit 3:
Energy Changes and Rates of Reaction

Thermodynamics

Definition & Importance

Understanding heat transfer properties is important
for building sufficient materials

All chemical reactions result in Heat Transfer

Different types of matter require different amount of
Heat Transfer to change the same temperature

Water is unusual - it can absorb/release a lot of
hear without the temperature drastically changing

Changes in a System

Endothermic Reaction

If the system absorbs energy and the
surroundings lose energy

Exothermic Reaction

If the system releases energy and the
surroundings gain energy

Calorimetry

The measure of heat change due to
a chemical reaction

A calorimeter is used to perform this task, while
a Bomb Calorimeter is a high tech version of this

Important Terms

Heat

The transfer of energy due to contact

Temperature

The measure of internal energy of an object
due to particle motion (kinetic energy)

First Law of Thermodynamics

The total amount of energy in the
universe is constant

Energy cannot be created or destroyed

Calculations

Specific Heat Capacity (c)

The amount of heat transfer required to
change the temperature of 1 gram of a substance
by 1 degree Celsius/Kelvin

c = specific heat capacity (in Joules/Grams*Celsius)
q = quantity of heat transferred (in Joules)
m = mass (in grams)
delta T = temperature (in degrees Celsius)

c = q/mT
q=mcT

Heat Capacity

The amount of heat transfer required to
raise the temperature of a sample by 1 degree
Celsius/Kelvin

c = heat capacity (in Joules/Celsius)
q = quantity of heat transferred (in Joules)
delta T = temperature (in degrees Celsius)

c = q/T
q= cT

Enthalpy

Total kinetic and potential energy of a
system at a constant pressure

Molar Enthalpy

The change in enthalpy on a per-mole basis

Enthalpy of Formation (delta Hf)

The amount of heat absorbed/released when 1 mole of the substance is formed at standard temperature (25 degrees Celsius) from its elements in their standard states

The (delta Hf) for elements in their standard states are ZERO

delta H = sum of (delta Hf of products) - sum of (delta Hf of reactants)

Hess's Law of Summation

The "H" is the same as the sum of the values
of the "H" for each individual step

Rules

#1. If all the coefficients of an equation are multiplied/ divided by a common factor, the "H" must be changed likewise

#2. When a reaction is reversed, the sign of "H" must also be reversed

#3. When cancelling compounds for Hess's Law, the state of the compounds is really important

Steps

#1. Number each given thermochemical equation

#2. Arrange the equations so that your desired reactants are on the LEFT side and the products are on the RIGHT side

#3. Multiply the equations by factors such that they match the desired equation (multiply the enthalpies by the same factor)

#4. Add the equations and enthalpies together (cancel out any repeating chemicals)

Reaction Rates

Rate of Reaction

How quickly reactants disappear to form products

Chemical reactions indicate the overall changes that is observed

Rate Theories

Collisions

Transition State

Catalysts

Reactions that require a GREATER number of particles to collide at the same time will DECREASE the chances of a successful reaction to occur

Factors Affecting Reaction Rates

#1. Chemical Nature

Precious metals are discovered first

Not very reactive

Alkali metals later founds (first found only in compounds)

Due to high reactivity

#2. Surface Area

Increase in surface area = Increase in the reaction rate

Heterogenous Reaction

The reactants are in different phases/states

Homogeneous Reaction

All the reactants are in the same phase/state

#3. Concentration

Higher the concentration = Higher the reaction rate

More chemicals result in more particles which can participate in the reaction

#4. Temperature

Higher the temperature = Higher the reaction rate

Increased temperature is due to increased particle motion

Greater the motion of a particle, the greater the chance it will encounter another reactant

#5. Catalyst

A compound that increases the rate of a chemical reaction without being consumed in the actual reaction

It lowers the amount of required energy for the reaction to occur

Also known as Activation Energy

Calculations

Rate Equation

rate = delta concentration / delta time

rate = K {A}^x {B}^y

Rate is measured in moles per litre per second (mol/L*s)

In the sense where A+B = C + D

The x and y values are determined by the actual experiment

K = the rate constant, determined by the reaction and the conditions the experiment was conducted in

Exponents of the Rate Law are NOT related to the coefficients, of the balanced chemical equation in any way

Rate Law Exponents

Exponents are related to the "Order of the Reaction"

Order of the Reaction

Experimentally determined by changing 1 reactant at a time and looking at how the reaction rate changes

The sum of the rate law exponents

Overall order = m

Unit of K = L ^(m-1) / mol ^(m-1) *s

Reaction Mechanisms

Occur over a set of stpes

Each step of the mechanism is known as "Elementary Process"

The overall reaction can only be as fast as the slowest elementary process (AKA Rate-Determining-Step)

Predicting Mechanisms

Only best guesses at the behaviour of molecules

Steps

#1. Each step must be elementary

Involve no more than 3 reactant molecules

#2. The slowest step (Rate-Determining-Step) must be consistent with the rate equation

#3. All the elementary steps must add up to the overall equation

Unit 5:
Electrochemistry

Net Ionic Equations

Any aqueous compounds completely ionize in water

Solids, liquids, or gases do NOT dissociate (completely ionize) in water

Steps

#1. Chemical equation

Normal equation written in the chemical names

#2. Turn the chemical equation into an Ionic equation

All the aqueous are separated into its component ions

#3. Complete the Net Ionic Equation

All the spectator ions (those which are the same on both sides of the equation) cancel out, and the final equation is written from the remaining components

At the end, the equation can be confirmed by counting all the atoms and charges on each side and making sure they are balanced

Oxidation and Reduction

Historical Definitions

Oxidation

Reaction of substances with oxygen (combustion OR corrosion)

Reduction

From metallurgy - producing metals from their compounds

Modern Definitions

Oxidation

The atom/ion that LOSES electrons

Reduction

The atom/ion that GAINS electrons

LEO the lion says GER

Loss of Electrons - Oxidation

Gain of Electrons - Reduction

Both reactions happen SIMULTANEOUSLY - referred to as REDOX reactions

Agents: One reactant causing a change in the other reactant

Oxidizing Agent

The reactant which causes the other to be OXIDIZED

Therefore, this reactant is reduced

Reducing Agent

The reactant which causes the other to be REDUCED

Therefore, this reactant is OXIDIZED

Half-Reactions

To monitor the transfer of electrons, represent each reaction seperately

A balanced chemical equation that shows the number of electrons involved

Oxidation Numbers

Used to keep track of electrons during reactions

An arbitrary system based on charge of ions and EN

Rules

#1. A pure element has an Ox. Number of ZERO

#2. The Ox. Number of an element in a monotomic ion equals the charge of the ion

#3. The Ox. Number of Hydrogen in its compounds is +1 except in metal hydrides, where the Ox. Number of Hydrogen is -1

#4. The Ox. Number of Oxygen in its compounds is usually -2, but there are certain exceptions

#5. In covalent compounds that do not contain "H" or "O" the more EN element is assigned an Ox. Number that equals the negative charge it usually has in its ionic compounds

#6. The sum of the Ox. Numbers of all the elements in a compound is ZERO

#7. The sum of all the Ox. Numbers of all the elements in a polyatomic ion equals the charge on the ion

Steps

#1. Assign common Ox. Numbers

Use the Periodic Table of Elements

#2. The total Ox. Numbers of a molecule/ion is the value of the charge of the molecule/ion

For neutral compounds, the Ox. Number of all the atoms must add up to ZERO

#3. The unknown Ox. Numbers can be assigned algebraically and be solved using variables

Unit 4:
Chemical Systems and Equilibrium

Equilibrium

Occurs when opposing changes (FWD and RVS reactions) are occurring simultaneously at the same rate

Theoretically, all reactions are reversible

Reactions wrtiten L-R are FWD reactions

Reactions written R-L are RVS reactions

Represented by an "equal sign" which has an arrow
pointing the opposite direction on each line

Both reactions occur until the concentrations of
reactants AND products undergo no further change

Keep in Mind

The reactant and product concentrations are CONSTANT, not equal

FWD reactions occur rapidly during the start, then slows down as reactant concentrations decrease

RVS reactions occur slowly during the start, then speed up as product concentrations increase

Generally, an equilibrium is a state of balance

EQLBM constant K

Dependence

Doesn't depend on the Reaction Mechanism

Doesn't depend on the initial concentrations

Depends on Temperature

Depends on the EQLBM concentrations

Magnitude

Since products are divided by reactants in the expression:

A larger K value (>1)

FWD reaction is favoured

A smaller K value (<1)

RVS reaction is favoured

All substances present are being made and unmade at the same rate

The Reactants are ALWAYS on the LEFT, and the Products are ALWAYS on the RIGHT side of an equation

At a given temperature, the reactants and products with ALWYAS be in the SAME RATIO at EQLBM, no matter the starting point

Le Chatelier's Principle

When a system at EQLBM is stressed, the syste, works to restore EQLBM once again

Types of Stress

Changes in Concentration

System shifts to get back to the same ratio of reactants and products

Means that Keq is constant

Changes to the Temperature

When temperature is increased, Endothermic reactions are favoured

Consider HEAT as a Reactant in Endothermic Reactions

A + B + Heat = AB

If heat is added

Shift to Products side

If heat is removed

Shift to Reactants side

When temperature is decreased, Exothermic reactions are favoured

Consider HEAT as a Product in Exothermic Reactions

A + B = AB + Heat

If heat is added

Shifts to Reactants side

If heat is removed

Shifts to Products side

Changes to the Pressure

Only affects EQLBM systems with unequal moles of Gaseous reactants and products

States of Equilibrium

Static

One in which there is NO motion

Dynamic

One in which there is motion, despite there being no Net Change

Chemical (an example of Dynamic)

The state where the reaction vessel contains a
mixture of all the reactants and products

Different Systems

Homogeneous

An EQLBM system in which all the components are in the SAME physical state (e.g. gas, aqueous, solid, liquid)

Heterogeneous

An EQLBM system in which the components are in DIFFERENT physical states (e.g. gas, aqueous, solid, liquid)

Calculations

Concentration

{C} = solute / solution

EQLBM Constant Expression

AKA Law of Chemical EQLBM , AKA Law of Mass Action

K = {products} / {reactants}

let the EQLBM sign be represented as {=}

E.G. aA + bB {=} cC + dD

K = {C}^c + {D}^d / {A}^a + {B}^b

Steps

#1. Calculate the molar concentrations

#2. Write the EQLBM expression

#3. Substitute the respective values and solve

Changes in Substances

Adding reactants

Shift to products side

Rf > Rr

Removing reactants

Shift to reactants side

Rf < Rr

Adding products

Shift to reactants side

Removing products

Shift to products side

RICE Table

R = Reaction

I = Initial Amounts (in concentration)

C = Change = Use variables

E = EQLBM value = Initial Amount - Change Variable

Solved using Keq = {products} / {reactants}

Acids and Bases

FIRST

Arrhenius Theory

Acid

pH level below 7

A substance that contains HYDROGEN in its chemical formula

Ionizes in water to form HYDRONIUM ION

Base

pH level above 7

A substance that contains HYDROXIDE ION (OH-) in its chemical formula

NEW

Bronsted-Lowry Theory

Acid

A proton DONOR

Base

A proton Acceptor

Conjugate

Acid-Base Pair

A pair of 2 substances related by the gain/loss of protons

Conjugate Acid

The particle formed after a proton JOINS a BASE

Conjugate Base

The particle remaining after a proton LEAVES an ACID

Strong vs Weak

Acid

Strong

An acid that COMPLETELY dissociates into ions in water

Weak

An acid that has LIMITED dissociation in water

The conjugate acid of a Strong Base is a weak acid

Base

Strong

A base that COMPLETELY dissociates into ions in water

Weak

A base that has LIMITED dissociation in water

The conjugate base of a Strong Acid is a weak base

Percent Dissociation

Ratio of concentration comparing ionized acid/base at EQLBM to original concentration of acid/base

Steps

#1. Calculate the concentration using pH

#2. Setup the RICE table and add values

#3. Create and solve teh EQLBM expression

#4. Calculate the Percent Dissociation

% Dissociated = ( {}dissociated / {}initial ) *100%

Kw = (Ka)*(Kb)

Can be re-arranged to solve for missing K_ value

Ka = Acid-Dissociation Constant for the Ionization of an Acid

Ka > Kb = solution is Acidic

Kb = Base-Dissociation Constant for the Ionization of a Base

Ka < Kb = solution is Basic

Final Concentration Calculations

Steps

#1. Determine the MOLES of both, acid and base

#2. Find the difference between the two values and determin which one has excess

#3. Re-calculate the concentration of the acid/base using the excess moles and the combined volume

Miscellaneous

Amphiprotic Substances

A molecule/ion which can accept/donate a proton

Can act as an acid or a base

Titrations

A-B Titration Curve

A graph of pH of an acid/base solution vs Amount of ADDED acid/base

Equivalence Point

The point in a Titration where the amount of Acid present exactly equals and reacts with amount of Base present

Steps

#1. Calculate the moles of Acid/Base

#2. Determine the excess moles and new concentration

#3. Use Kw to calculate the concentration and find pH

Solubility (Qsp)

Molar Solubility

The amount (in moles) of solution in 1L of saturated solution

Precipitate

An insoluble product that forms from a reaction between 2 soluble ionic compounds

Vraj Rao - Mr. Z. Syed - SCH4U Grade 12 Chemistry

Unit 1: Structures and Properties of Matter

Unit 2:Organic Chemistry

Unit 3:Energy Changes and Rates of Reaction

Unit 5:Electrochemistry

Unit 4:Chemical Systems and Equilibrium

Unit 1:
Structures and Properties of Matter

Unit 2:
Organic Chemistry

Unit 3:
Energy Changes and Rates of Reaction

Unit 5:
Electrochemistry

Unit 4:
Chemical Systems and Equilibrium