UNIT 1: ATOMIC STRUCTURES
ATOMS: building blocks of molecules
ELEMENT: matter composted of one atom (e.g., gold, diatomic O2 gas)
MOLECULE: 2 or more atoms bonded together (e.g., O2)
COMPOUNDS: molecules of more than one type of atom (e.g., ammonia, water)
All atoms are composed of subatomic
particles (protons, neutrons, elections)
Smallest to largest = electron > proton >
neutron
PROTONS: (+), 1amu = 1 dalton, in nucleus
ELECTRON: (-), NO MASS; (-) = (+) on proton
(magnitude = 1.6 x 10 -19 coulombs; found in
electron cloud (surrounding nucleus)
NEUTRONS: neutral, 1amu, in nucleus
- Electron cloud has ATOMIC ORBITALS (shells) where electrons circle nucleus
o Opposite charges attract
- Electrons in orbitals closest to nucleus are lower in energy and most stable
- Electrons in orbitals further away from nucleus are higher in energy and less stable
VALENCE SHELL: outermost shell; responsible for interacting and reacting with other atom
electrons.
ATOMIC IDENTITY
MASS NUMBER (A): protons +
neutrons; top left corner
ATOMIC NUMBER (Z): number of
protons; gives element identity
*mass number – atomic number = # of
neutrons
ISOTOPES: same element but different number of neutrons (e.g., Carbon-12, Carbon-13); labels
and tracks atoms
ATOMIC WEIGHT: average mass of atom of an element (accounts for isotopes)
- E.g., carbon has an atomic weight of 12.01 amu
MOLECULAR WEIGHT: WEIGHT OF MOLECULE
- Atoms with the same element but different number of neutrons and electrons
IONS: atoms that lose or gain electrons and become charged in valence shell
ELEMENTAL STATE: default state; neutral; # protons = # electrons
CATION: + charged ion (electron poor, proton rich)
ANION: - charged ion, electron rich, proton poor)
-
LEFT SIDE: love to lose electrons to form + cations
RIGHT SIDE: love to gain electrons to form – anions
o *NOTE: noble gases are relatively stable
MONATOMIC IONS: type of anions; uses suffix -ide (ex. Hydride, oxide)
OXYANIONS: anions that contain O2; uses suffix
- _-ite,
- hypo__ite (1 less than ite)
- _-ate (for more O2 atoms)
- Per_ate
- Ex. Hypochlorite (ClO-), Chlorite (ClO-2) Chlorate (ClO-3), Perchlorate (ClO-4)
POLYATOMIC IONS: ions with more than 1 atom (ex. NH4, H3O+)
- Elements with 1+ cation are indicated with Roman Numeral or suffice, -ous and -ic
o Ex. Fe2+, Iron (II), ferrous ion; smaller charge
o Ex. Fe3+, Iron (III), ferric ion; larger charge
- Polyatomic ions with hydrogen atoms use hydrogen and dihydrogen before parent
name
o Ex. CO32- carbonate -> HCO3- hydrogen carbonate
o Ex. PO43- Phosphate -> H2PO4- dihydrogen phosphate
THE BOHR MODEL
1805 – Atomic theory was postulated by John Dalton in
- All matter consists of atoms that form elements and participate in chemical reactions
1829 – Cell Theory
1861 – Germ Theory
1892 – Viruses discovered
1913 – Bohr Model
- Neil Bohr and Ernest Rutherford introduced this model for structure of hydrogen atom
1943 – DNA Helix discovered
RECALL: electrons orbit atomic nucleus in spherical shells at varying distances from nucleus
- + charged protons in nucleus exert attractive forces on electrons
- The closer the electron to the nucleus, the greater the attractive pull making electrons
more stable and lower in energy
Electrons are ordered with increasing
distances from the nucleus
N = 1: closet orbital; lowest energy level;
ground state
For electron to jump from lower to higher
energy orbital, must absorb discrete amount
of energy in form of photon of light
- Electron becomes excited
For electron jumps from high energy orbital
to lower, it will emit a photon of light
Energy emitted or absorbed is in form of
ELECTROMAGNETIC RADIATION
ELECTROMAGNETIC RADIATION: radiation with both electric and magnetic fields
- Can include visible light, radio waves, gamma rays, X rays
Amount of energy associated with one of these photons of light is equal to Planck’s Constant, H
- E = hf
o h = Planck’s constant
o f = frequency of light
Frequency of wave
- f = c/λ
- can also be written E = hc/ λ
Rutherford and Bohr demonstrated that when electrons jump between energy levels, energy of
photon is emitted absorbed in discrete amounts called QUANTA
- photons cannot just have any amount of energy it wants; MUST BE EXACT VALUE so we
can calculate/predict it
- results in ATOMIC EMISSION SPECTRUM: pattern of distinct lines that each represent
quantized amount of energy emitted when electron falls to lower energy level
o ATOMIC ABSORPTION SPECTRUM: each line corresponds to quantized amount
of energy absorbed during increase in energy level
o unique for each element
RYDBERG FORMULA: amount of energy emitted/absorbed can be approximated using this
formula
- change in energy is equal to Planck’s constant
-
ΔE = hc/ λ = R (1/nni – 1/n2f)
R = 2.18 x 10-18 J
H = 6.626 x 10-34 J
-
ΔE = (2.18 x 10-18)(1/12-1/32)
= (2.18 x 10-18)(8/9)
= 1.4 x 10-18J
QUANTUM NUMBER
HEISENBUERG UNCERTAINTY PRINCIPLE: it’s impossible to know the exact location and exact
momentum of an electron
- Have some information about where electrons are and what they are doing
- Exist in ORBITALS: areas of space where given electron is likely to be found by mathematical
probability functions
o Each orbital is part of a specific subshell
o Each subshell is part of larger region called ELECTRON SHELL
QUANTUM NUMBER
Every subshell has a unique quantum number that give an address within the atom
According to PAULI EXCLUSION PRINCIPLE: no two electrons can have same 4 quantum numbers
1. PRINCIPLE QUANTUM NUMBER (N): describes the energy level of electrons (n=1,2,3,4) or the
shell it is in
the higher the principal quantum number, the farther
the shell is from the nucleus
- greater the energy associated with that energy
level
- increases down the periodic table
2. ANGULAR MOMENT / AZIMUTHAL QUANTUM NUMBER: describes the shape of the orbital
Tells us what subshell electron is located in within a shell,
where L can range 0 – n-1 for given principal quantum
number
L = 0 represents S subshell
L = 1 represents P subshell
L = 2 represents D subshell
L = 3 represents F subshell
Shape shows probability of finding an electron in space
3. MAGNETIC QUANTUM UNMBER (Ml): tells us spatial orientation of the orbital (-l to +l)
ex. Whether a particular orbital within P subshell is
parallel to X,Y,Z axis in space
Regardless of what subshell it is found in, each
orbital can hold only two electrons
Ml values range from -l to +l
S subshell has 1 orbital where Ml = 0 so one possible
space in orientation for this subshell
P subshell holds 3 orbitals where Ml = -1, Ml = 0 and
Ml = +1
D subshell holds 5 orbitals
F subshells hold 7 orbitals
*If any orbital can only contain at most, 2 electrons, S subshell must be able to hold up to 2 electrons
- P subshell can hold up to 6
- D subshell can hold up to 10
- F subshell can hold up to 14
4. SPIN QUANTUM NUMBER (ms): spin orientation of these electrons
Two electrons in orbital differ in their spin, a form
of angular momentum
Can have negative half (-1/2) or a positive half
(+1/2)
PAIRED ELECTRONS: two electrons in the same
orbital; must have opposite spins
SHELL: tells orbit path electron follows around the nucleus; described by principal quantum number
Each shell is comprised of SUBSHELLS (s,p,d,f subshells) which have unique shape indicating where we
are most likely to find electrons
ORBITAL: subshells are consisting of orbitals which all have unique spatial orientation and can hold up to
2 electrons
EXAMPLE: how many electrons can exist at energy level n=4?
ELECTRON CONFIGURATION
No electron can share exact same 4 quantum numbers
Electrons like to be close to nucleus as possible where energy
is low and most stable
Can predict electrons will fill lowest orbital energy first before
filling higher energy orbitals which generates most stable
electron configuration
AUFBAU PRINCIPLE: electrons fill orbitals in order of lowest to
highest energy; used to find electron configuration of
individual atoms
Subshells increase in energy with increasing principal
quantum numbers
- Also increase in energy with increasing angular
momentum quantum numbers (S to P ex.)
SHORTCUT TO MEMORIZE SUBSHELLS IN INCREASING ENERGY
1. Write numbers 1 to 5/6 down column in page
2. Fill subshells from left to right
3. Fill in order by drawing arrows diagonaly from top right to bottom left corner
o ex. 1S is filled first, then 2S, 2P, 3S, 3P etc.
ANOTHER SHORTCUT
Divide periodic table of elements into blocks by angular momentum number which helps determine
subshell in which valence electron will be located for given element
- 1S block begins in first row
- 2S block begins in second row
- 3D block begins at fourth row
- 4F block begins at row 6
1. Abbreviate electron configuration by placing last element of prior row in brackets to represent
its electron configuration
2. Add configuration of valence electron from periodic table
a. Ex. For chlorine, we’d place neon in brackets to represent its electron configuration and
then electron configuration for chlorine’s valence electrons
b. Chlorine has 7 valence electrons ([Ne]3s23p5) as its in 7th column of periodic table
c. Each box that proceeds chlorine is a placeholder for one of its valence electrons
UNIT 2: PERIODIC STRUCTURES AND TRENDS
2.1 STRUCTURE OF PERIODIC TABLE
PERIODS: rows of elements
- Number of valence electrons increase until valence shell is full
- Ex. Second period: Li = 1 valence, Be = 2, B =3, C=4 …
GROUPS/FAMILIES: columns of elements
- Elements have the same number of valence electrons
- Share physical and chemical properties
- Increasing electrons
o Elements with fewer electrons are placed on top, more on the bottom
- Top of periodic table contain elements where valence electrons are at lower energy levels
LANTHANIDE SERIES: pink
ANTINIDE SERIES: purple
Elements are divided into 3 main categories
1. Metals
o Far left of periodic table and throughout middle block
o Good metal and electrical conductors
o Solid under standard conditions
o Shiny (lustrous) , ductile, malleability
 Can be easily stretched into wires or other shapes
2. Nonmetals
o Not lustrous
o Poor electrical conductors
o Far right of periodic table
3. Metalloids
o Share traits of metals and non-metals
o Ex. Boron, silicon
o Brittle
o Poor to decent electrical conductors
o Found in narrow, staircase row between transition metals and nonmetals on the right
GROUPS
- IUPAC Naming System
o International Union of Pure and Applied Chemistry
o Denoted by roman numeral naming system of letters A and B
o Group numbers correspond to number of valence electrons in neutral atom (exception =
transition metals)
o NOW: numbers them from 1-18 moving left to right
- GROUP 1: Alkali Metals
o Exception of H (nonmetal)
o All group 1 elements have only one valence electron
o Highly reactive
 Ex. Na, K are so reactive they explode when added from solid form into water
o Readily donate their valence electron and form cations with +1 charge or oxidation
number
-
-
-
-
-
-
-
-
GROUP 2: Alkali Earth Metals
o Have 2 valence electrons they donate to form +2 cations
o Most common = calcium and magnesium
o Metallic and reactive
GROUP 3-12: Transition Metals
o Have gold and platinum which are valued for luster and unique properties
o Hard, durable metals that readily conducts electricity
o Take on vivid colours due to electronic transitions between d orbitals
o Number of electrons it can lose varies
o Most have multiple oxidation states
 Ex. Fe can have oxidation numbers of +2 or +3 but -2 to +6 possible
GROUP 13: Metals
o Has metalloid boron but rest are metals
o Aluminum is most common
o 3 valence electrons allowing them to form +3 cations
o Have ability to form multiple oxidation states
GROUP 14: Carbon family (nonmetal)
o Silicon and Germanium are metalloids
o Tin and Lead are metals
o Carbon has tendency to form 4 bonds
o Carbon can take many forms (e.g., diamond = hard solid while also being a gray solid
material known as graphite)
o Can form oxides (e.g., carbon dioxide, silicon dioxide)
o 4 valence electrons
GROUP 15: Nitrogen Family
o 5 valence electrons
o Most common are nitrogen and phosphorus (nonmetals)
 Arsenic and antimony are metalloids
 Bismuth is metal
GROUP 16: Chalcogens
o Include Oxygen, Sulfur
o Oxygen, Sulfur and Selenium are nonmetals
 Tellurium and Polonium are metalloids
o Need 2 more electrons to be filled
 Elements react with other atoms that causes them to gain 2 electrons
 Forming anions with -2 charge
GROUP 17: Halogens
o Nonmetal elements
o Need one more electron to complete valence shell
o Highly reactive
o Common pattern = they accept an electron from group 1 element to obtain -1 charge
o Tendency to fill 1 electron shell explains fact that in natural state, found in diatomic,
covalently bonded molecules (e.g. Cl2, Br2, I2)
GROUP 18: Noble Gases
o Non-metallic character and unreactive due to full valence shell
o Low boiling points causing them to exist in gaseous state under standard conditions
PERIODIC TRENDS
- Effective nuclear charge (Zeff): attractive force of + charged nucleus on atom’s valence
electrons (NOT NUMBER OF PROTONS)
o Amount of attractive force being induced on outermost electron
o As we move across row/table, effective nuclear charge will increase because
atomic number is increasing (left to right), along with # protons in nucleus and
the amount of charge valence electrons are exposed to
o Effective nuclear charge increases as it moves from left to right across a period
o More protons in nucleus and less non-valence electrons increase Zeff
- If we move down column/group, there are increasing electron shells located between
nucleus and valence electrons – shield valence elections from impact of positive charge
of nucleus.
o Effective nuclear charge decreases as we move down a group
ATOMC RADIUS AND IONIC RADIUS
- Show highly predictable trends that are derived from patterns in effective nuclear
charge
- If effective nuclear charge is stronger (protons of nucleus can pull valence electrons
more tightly) radius is smaller
- Atomic radius: radius of neutral atoms or atoms in elemental state
o moving across period  effective nuclear charge increases (additional protons
pull electrons closer, decreasing atomic radius)
o going down groups  Zeff decreases – electrons held in place more loosely by
attractive force, increasing atomic radius
o atomic radius shows inverse relationship with effective nuclear charge
- ionic radius: radius of ions or charged species
o able to predict the radius of cations (positive ions) and anions (negative ions)
relate to radius of uncharged counterparts
- electrostatic repulsion among electrons
o like charges repel – when we add electrons to form anion, electrons will want to
get away from eachother more, increasing radius of ion relative to neutral atom
o removing electrons to form cation will make radius smaller that neutral atom
- ionization energy: energy needed to remove one valence electron from a neutral atom
in gaseous state
o positive quantity as energy is needed to pull electron away from nucleus
o first ionization energy is lower than the second (less energy is required) because
removing an electron reduces electrostatics repulsion (gives boost to remove
first electron easy to remove as reflected by lower first ionization energy)
- it is hard to remove electrons from a stable configurations
o ex. Na has 1 valence electron and 3s orbital – removing that valence electron
produces configuration with full 2p orbital (like noble gases – highly stable) so
second ionization energy is Na is dramatically higher (10x higher than 1 st
ionization energy)
LESSON 3.1: BONDING
ATOMS, IONS, MOLECULES, COMPOUNDS
Atomic Theory
- all matter consists of indivisible atoms (radio active decay is exception)
- atoms of the same chemicals are identical.
- compounds consist of atoms of more than one element connected (ex. H2O =
compound)
- molecules: any structure composed of multiple atoms (ex. H2)
- atoms: consist of protons, neutrons, electrons
o electrons = “orbiting planets”
o protons + neutrons = the sun
o mostly composed of empty space
o protons and neutrons larger than electrons (x2000)
o protons and electrons are equivalent in charge (1.6 x10-19 coulombs) while
electrons are -1.6x10-19 coulombs
o neutrons are neutral and atoms in their natural or elemental state occur with a
neutral charge (same number of protons and electrons)
IONS
-
-
Ion: any charged species
caTion: + charged ion
anion: negatively charged species
MEMORIZE THESE IONS
o Polyatomic ions (charged species consisting of multiple atoms bonded together:
1. Nitrate (NO3-)
2. Nitrite (NO2-)
3. Carbonate (CO32-)
4. Bicarbonate (HCO3-)
5. Perchlorate (CIO4-)
6. Chlorate (ClO3-)
7. Sulfate (SO42-)
8. Sulfite (SO32-)
9. Hydroxide (OH-)
10. Chromate (CrO42-)
11. Cyanide (CN-)
12. Permanganate (MnO4-)
13. Acetate (C2H3O2-)
Binary Compound: compound made up of two elements
Ionic Compound: contains a metal and a nonmetal (KCl potassium chloride)
o List Cation first and Anion second (K  cation, Cl  anion)
o Metal retains normal name followed by nonmetal with suffix -ide
o No prefixes used when naming ionic compounds.
 Ex. CaF2 will be calcium fluoride, not calcium difluoride
-
o Naming ionic compounds reflects characteristic charges – depending on where
each element is found on the table, can predict whether it forms cations or
anions when ionized.
 Elements found on left and right side = can predict how many electrons
will be gained and lost
 Transition metal: can form different ions in different situations
o Transition Metals: oxidation number specified with Roman numerals (ex. Fe (III),
Fe (II)
 tells us how many electrons transition metal has lost which provides info
to predict how many atoms of corresponding metal partnered
 ex. ZnO2 is Zinc (IV) Oxide whereas ZnO is zinc (II) oxide
Molecular Compounds: contain two nonmetals (NO2 nitrogen dioxide)
o Elements arranged in order of electronegativity
 EXCEPTION: carbon first, hydrogen (2.1 EN) is written after nitrogen (3.0
EN)
o “-ide” is added to end of the name of second element
o Prefixes used to indicate how many atoms of each element are present
 1 = mono, 2 = di, 3= tri, 4= tetra, 5 = penta.
 Mono left out if there is one atom in first element
 CO is carbon monoxide, not monocarbon monoxide
BONDING/INTRAMOLECULAR ATTRACTION
ATOMS
- Most atoms appear as a molecule as many are unstable on their own as they have
incomplete valence shell (except noble gases)
o Noble gases are chemically inert, don’t bond with anything as valence shells are
full
- Atoms can achieve more stable state by coming together
-
Intramolecular forces: forces that hold atoms together in molecules – there are 3 types
1. Covalent bonds: formed between two nonmetals by sharing valence electrons
a. Occurs because non metals have relatively similar electronegativity values
b. Nonpolar Covalent Bonds: bonding electrons share equally between two
atoms. No changes on the atoms.
c. Polar Covalent Bond: bonding electrons shared unequally between two
atoms. Partial charges on atoms
i. because electrons are shared unequally in polar covalent bonds, a
stable dipole forms across the bond (there is electron surplus and
partial negative charge on more electronegative atom, and an
electron deficiency and partial positive charge on less
electronegative atom)
ii. the dipole formed across the polar covalent bond is a dipole
moment: vector formed by charge differential
iii. BE CAREFUL: as molecules can contain multiple covalent bonds and
overall dipole of molecule is by adding dipoles of all the bonds –
leading to traps (CaCl4) where they cancel out
d. Pauling Scale:
i. F is most electronegative element with score of 4.0 while Cs has value
of 0.79, H is 2.2
ii. Electronegativity differences <0.5 correspond to nonpolar covalent
bonds
iii. Electronegativity differences between 0.5 and 1.7 correspond to polar
covalent bonds
2. Ionic bonds: formed when electronegativity difference between atoms exceeds 1.7
a. Form between metals and non-metals = stronger than covalent bonds
b. Electrons completely transferred from metal to non metal
c. metal to become a + charged cation while nonmetal becomes – charged
anion
d. ex. MgCl2 has 2 ionic bonds, Mg is cation with oxidation # of +2, so it gives up
both its electrons while Cl is anion with oxidation # of -1, so it receives 1
electron from Mg
3. Metallic bonds: result from metal atoms joining together where electrons become
delocalize and form sea of electrons that are free to move throughout solid
a. Sea of electrons accounts for many properties of metals including conducting
heat and electricity
NOTE: C-H = nonpolar
N-H = moderately polar
O-H, F-H are highly polar
C-H = slightly polar
C-O = meaningfully polar (not as much as OH)
INTERMOLECULAR FORCES
- Intramolecular forces: connect atoms within a molecule.
- Intermolecular forces: interaction between molecules
o Weaker than intramolecular forces
- Water has 2 covalent bonds (intramolecular polar covalent
bonds ) which represent intramolecular forces that represent the
joining of the oxygen and hydrogen atoms
- Between the water molecule, there is another attractive force
between oxygen atom of one molecule and the hydrogen atom of
another (intermolecular force between partial negative charge of
oxygen and partiail positive charge of hydrogen
- NOTE: partial charges exist because bond between O and H is a
polar covalent bond
INTERMOLECULAR INTERACTIONS
- Intramolecular forces create polar molecules that can be connected to other polar molecules via
intermolecular forces.
- Larger dipole moments facilitate stronger intermolecular forces
- Intermolecular forces are divided into 5 types in order of increasing strength:
1. London Dispersion Forces (weakest)
a. Can occur between any molecule regardless of how nonpolar they are – arise from
random alignment of electrons
b. Weak force will exist between positive and negative areas of charge
c. Technically present between any pair of molecules (may be so small that there is no
difference)
d. Occurs when temporary dipoles arise by chance
e. Larger structure = larger London Dispersion Force
2. Dipole-Dipole Interactions
i. Occur between stable dipoles (polar molecules)
ii. Attractive force occurs between positive dipole of one polar molecule and
the negative dipole of another molecule
iii. Ex. Acetone – O has – charge and partial + on central carbon, so O atom of
one acetone molecule will be attracted to central carbon of another
3.
o
o
o
o
o
o
Hydrogen Bonds
Occurs when a H attached to N,O or F is attracted to lone pair
on a N,O,F
Ex. H2O and H-F – hydrogen in H-F has partial positive charge
and attracted to lone pair present in O in H2O
H2O’s hydrogen can also hydrogen bond with F in H-F
ONLY occurs with N,O,F because these atoms are very
electronegative and generate strong dipoles with H
The greater the difference in electronegativity between 2 atoms
in a molecule, more polar the bond within molecule will be
Strong stable dipoles lead to strong intermolecular interactions
4. Ion Dipole Forces
a. Occur between ions and molecules with a dipole – ions have full charges (extra
polar)
b. Occur between two molecules with a full charge.
c. Intermolecular force
5. Ionic Interactions (strongest)
RECAP:
- Intermolecular forces allow us to make predictions about how substances behave
- Melting and boiling points
o Reflect thermal energy required to break down attractive forces of a substance
- The stronger intermolecular forces a compound has, the higher its melting and boiling points
LESSON: CHEMICAL REACTIONS:
TYPES OF REACTIONS
- Reactions: process where chemical species are converted into one or more chemically different
species
o Reactants  products
- Not all noticeable changes to chemical compounds are reactions.
o Ex. Phase changes (ice  water) is not a reaction as no bonds were broken/formed.
- Bonds need to be broken or formed.
- Atoms involved in chemical reactions remain largely unchanged, except for rearrangement of
valence electrons.
o Ex. If you start reactant only containing carbon, hydrogen and oxygen, your product will
also only contain carbon, oxygen and hydrogen
o EXCEPTION: radioactive decay
BASIC REACTION TYPES (7)
BASIC REACTION
RXN
TYPE
SYNTHESIS
A + B C
DECOMPOSITION
A B + C
DEFINITION
EXAMPLE
Two or more reactants
combine to synthesize a
single product
N2(g) + 3H2(g)  2NH3(g)
Single reactant breaks down
into multiple products
H2CO2  H2O + CO2
endothermic energy must be
invested to break chemical
bonds
∆H>0
SINGLE
DISPLACEMENT
AB + C AC + B
One element or group of
atoms replace another
within a molecule
Often involve an ionic
compound with either
cationic or anionic part of
ionic compound is replaced
by chemically similar cation
or anion compound
DOUBLE
DISPLACEMENT
AB+CD  AD + BC
NEUTRALIZATION Acid + Base  H2O
+ Salt
Two iconic compounds are
exchanging their anions or
cations to form two new
compounds
Subtype of double
displacement reaction
FeSO4(aq) + 2Na(s) 
Na2SO4(aq) + Fe (s)
Cl2(g) + 2NAI(aq)  2NaCl (aq)
+ I2(s)
NaBr(aq) + AgNO3(aq) 
AgBr(s) + NaNo3(aq)
NOTE: AgBr is insoluble
precipitate
HCl (aq) + KOH (aq)  H2O (I) +
KCl (aq)
Acid contains hydrogen, Base
contains hydroxide
Can result in solution with
neutral pH (7) but not always
case
COMBUSTION
Compound reacts (burns) in
presence of oxygen
C3H8 + 5O2  3CO2 + 4H2O
Most rxns include
hydrocarbon reacting with
oxygen to form CO and
Water (ex. Combustion of
propane or alcohol)
Combustible organic
molecules contain C, H, or O
Highly exothermic. ∆H<0
OXIDATION
REDUCTION
Characterized by electron
transfer; species that gives
up electron is oxidized,
species that gains electron is
reduced
OXIDATION REDUCTION EXAMPLE
REACTIVITY SERIES (ACTIVITY SERIES)
- Lists relative reactivity of different metals
- Most to Least = lithium, sodium, iron, gold
BALACING CHEMICAL EQUATIONS
- Conservation of mass: mass in an isolated system is neither created or destroyed
- Any atom appearing in the reactants must be accounted for in the products
-
Lack of coefficient in front of any of the reactants or molecules suggest that it’s most likely
unbalanced.
STEPS FOR BALANCING
- Group together polyatomic ions that are the same on both sides (treat NH4 and OH as individual
units)
- Look for elements or groups that appear in only one species on each side. (Fe, Cl, NH4, OH)
- Save elements that appear by themselves for last
- Important to write the process out and work systematically
- Law of conservation applies to not only mass but charge as well – net charge of each side of a
chemical of an equation must be balanced out
EXAMPLE FOR BALANCING EQUATIONS
LESSON: LEWIS STRUCTURE AND VSEPR
OCTET RULE
-
-
OCTET RULE: atoms prefer to have 8 electrons in their valence shell – only consider P (6-) and S
(2) e-) electrons involved
EXCEPTIONS TO RULE:
o Hydrogen – has single valence electron subshell (1s orbital) and can only hold 2
electrons so it’s stable with 2 electrons.
o Helium – Max: 2 valence electrons
o Lithium – 1 valence electrons but can lose 1 electron (LiH) = Max: 2 valence electrons
o Beryllium – 2 valence electrons but can gain 2 electrons (BeH2) = Max: 4 valence
electrons
o Boron: 3 valence electrons but can gain 3 electrons = Max: 6 valence electrons
EXPANDED OCTET: ability for valence shell to hold more than 8 electrons
HYPERVALENT MOLECULES: molecules in central atom has more than 8 valence electrons
Ex. Phosphorous Pentafluoride (PF5) – central phosphorus atom has 10 valence electrons
Ex. sulfur hexafluoride – where the sulfur atom has 12 valence electrons
OCTET RULE EXCEPTION
- Odd Numbered Electrons
o Ex. nitrous oxide contains 5 valence electrons from nitrogen and 6 from oxygen (total of
11 valence electrons) and having 8 around each atom is impossible.
 Also has 1 unpaired electron
 Known as free radical – highly unstable
LEWIS STRUCTURE
- Dots: represent number of valence electrons
- Lines: represent 2 shared electrons connecting two atoms – double/triple bonds represented by
2/3 lines and - or + denote presence of negative charge due to lack or excess of electrons
- two single bonds between O and H and 2 lone pairs on oxygen
- Each single line corresponds to 2 electrons – as they are shared, can satisfy
octet rule
DRAWING LEWIS STRUCTURES
1. Draw molecules in proper positioning – need to determine what atom is central atom and
arrange other atoms around it (H is never central atom)
a. Atom with lower electronegativity is usually always central.
2. Determine how many valence electrons must be present in entire molecule
3. Atoms must be connected by at least on bond
a. NOTE: if central atom does not have complete octet, start erasing lone pairs with
peripheral atoms and replace each with multiple bond to central atom until all atoms
have proper octets
b.
RESONANCE STRUCTURES
resonance structures must have same
positioning and connectivity of atoms, but differ in
distribution of electrons across the molecule.
Example: nitrate anion, NO3- which has 3 valid
lewis structure bonds
-
When trying to figure out if two structures are resonance forms of eachother, MUST REMEMBER
o All structures have the same number of electrons – can differ depending on if its bonds
or lone pairs
o When converting resonance pairs into another, shift lone pairs or pi electrons to
adjacent atoms as opposed to atoms more distant on molecule
o Atoms do not change position in resonance forms
most stable resonance structure contributing most to its
behaviour OR can say resonance structures are ways to
capture aspects of a molecule’s structure and present them
in visual way
RESONANCE STRUCTURE OF O3 (OZONE)
-
-
Ozone has 2 resonance structures, each are equally stable therefore, contributing equally to
ozone’s actual structure
Resonance hybrid of ozone contains 2 partial double bonds with even distribution of electrons
Bond order (more than 2 atoms in molecule) =
The even distribution of electrons in hybrid structure of ozone reflects ELECTRON
DELOCALIZATION which explains why resonance increases the stability of molecules.
FORMAL CHARGES
central oxygen atom must have one lone pair to have
full octet
When determining number of valence electrons, only count one of the
electrons in each bond
-
-
Formal Charge: disparity between the number of electrons an atom should have and the
number it actually has
o Ex. for Ozone, the presence of 5 electrons reflects deficiency of electrons, yielding
formal charge of +1
o Helps evaluate stability of resonance structures
o Better resonance structure minimize formal charge
Formal Charge (N)= valence electron – lone pair electrons – ½ bonding electrons
o = should – actually has
Example: NH3
1. Write out Lewis structure diagram
2. Use Formula for Formal Charge
a. Formal Charge = valence electron – lone pairs – ½ bonding
electron
b. Formal charge (N) = 5-2 -6/3 = 0
NOTE: atoms that have fewer electrons than the periodic table have
positive formal charges
-
-
Atoms that have more electrons than the periodic table have negative formal charges
Formal charges ≠ actual charges
o B/c formal charge calculations assume that electrons are evenly shared between atoms
that are covalently bonded with eachother
o Highly electronegative atoms disproportionately pull electron density towards
themselves, which can lead to errors in predicted and actual charge distributions
Ex. H2O – formal charges on all water atoms are 0 and molecule has net 0 charge but O is
electronegative and pulls electrons towards itself
o Oxygen has partial – charge and hydrogen has partial + charge
ORBIT HYBRIDIZATION AND VSEPR THEORY
- Electron’s move in specific patterns at varying distances from the nucleus
ELECTRON SHELLS: different differences – regions of
high electron density
Each shell contains at least one atomic orbital
Atomic orbitals depend on the angular momentum
quantum number and refer to patterns in which
electrons move
MOST COMMON ATOMIC ORBITAL  found in s subshell
(spherical) and p subshell (two lobes and central node) – d and f
exist but geometry becomes complex.
ORBITAL BONDS
- When electrons join bonds together to generate molecule,
atomic orbitals combine to form molecular orbital
- Sigma (σ) bond is created when single bond is formed between
2 atoms – large area overlapping electron density compared to pi
bonds
- Pi (π)Bonds: occurs between 2 parallel p orbitals and are
weaker than sigma bonds
-
Double bonds = 1 σ + 1 π
Triple Bonds = 2 σ + 1 π
NOTE: in atoms that form four single bonds, orbitals will combine or hybridize to produce 4 equal sp3
orbitals
HYBRIDIZATION
- Between 1s and 2p orbitals may also occur
o S+p+p  3 sp2
o
S+p  2sp
- To identify hybridization, need to determine the number of regions of electron
density around the atom.
o
Region of electron density: bond, single, double, or triple or lone pair of
electrons
- Two regions yield a hybridization of sp (ex. cases of triple bonds or central
atoms with 2 double bonds
EXAMPLE: CARBON DIOXIDE
-
central C atom bonded to to other atoms, O, with 0 lone pairs
There are only 2 regions of electron density
Carbon atom ins CO2 has sp hybridization
VALENCE SHELL ELECTRON PAIR REPULSION THEORY: uses Lewis struct. And electronic relationships to
determine shapes of molecules assuming that the distance between eletron rich regions will be
maximized due to mutual repulsion of electrons (like charges repel)
ELETRONIC GEOMETRY: takes into account lone pairs and bonded atoms when producing shape
MOLECULAR SHAPE: considers only bonded atom even though these bonds are repelled by any lone
pairs present in central atom
- Ex. NH3  central nitrogen atom is attached to 3 bonded atoms and one lone pair, giving it 4
regions of electron density – electric geometry is tetrahedral and molecular shape is trigonal
pyramidal (lone pairs push bonds into pyramid like shape)
MOLECULAR SHAPES
- Tetrahedral molecular shapes: central atom makes 4 different bonds to other
substituents, occurs in carbon containing molecules with no double/triple bonds
such as methane (CH4)
o
To maximize separation of atoms bound on central atom, angle on each
bond is 109.5 degrees
LESSON: PHASE AND PHASE CHANGES
PHASES OF MATTER AND THEIR CHARACTERISTICS
SOLIDS
- A rigid shape due to packing and fixed volume, and particles don’t flow and move past one
another.
- Two types of solids:
o Crystalline Solids: packed in ordered manner with repeating pattern. ex. rock salt where
each sodium ion is positioned equidistant to 6 neighbouring chloride ions
o Amorphous Solids: have no order to their particles and are not as rigid as crystalline but
have strong material. ex. glass, gels, polymers (rubber, plastics)
- NOTE: 4 types of crystalline solids
o Ionic – made of ionic compounds (NaCl). Can see from crystal structure how ions are
specifically ordered to optimize electrostatic interactions.
o Molecular – molecules held together by non-ionic intermolecular forces (dipole-dipole,
London dispersions). Ex. Ice
o Covalent Network – certain molecules held together by covalent bonds. Ex. diamonds
have carbon atoms covalently bonded in regular pattern.
o Metallic Solids – contains metal atoms held together by metallic bonds. Ex. Gold, where
delocalized valence electrons in metallic solids give metal ability to conduct heat and
electricity.
LIQUIDS
- Have no fixed shapes, hard to compress and flow readily.
- Viscosity: resistance of liquid to flow (ex. molasses vs water)
- Surface Tension: amount of energy needed to increase the surface area of a liquid. (Tension
created at liquid surface by intermolecular attractions btwn particles)
o Interactions are stronger (energetically favourable) than corresponding intermolecular
attractions between liquids and surroundings.
o Results in liquid to take up the least space possible and minimizes surface area.
- Adhesive Forces: attractive forces between molecules of 2 different substances.
- Cohesive Forces: attractive forces between molecules of the same substance.
o These 2 types of forces interact to produce macroscopic effect of capillary action
o Capillary Action: movement of liquid up the sides of narrow tubes against force of
gravity.
 Occurs due to adhesion of liquid molecules to sides of the tube, which then pull
lower liquid molecules upwards by cohesion.
GASES
- Flow and have no fixed shape, and have any volume (expandable and compressible)
- Particles are free moving and are so spaced apart that individual particles have little interactions
with other particles.
- In a container of gas, each gas particle is moving randomly, bouncing off other gas molecules
and wall of container.
-
PHASE TRANSITIONS
- Solid  liquid = melting/fusion
- Liquid  gas = vaporization
- gas  liquid = condensation
- liquid  solid = freezing
- solid  gas = sublimination (ex. CO2 or dry ice transition to gas phase)
- gas  solid = deposition
GIBBS FREE ENERGY
- ΔG<0 = spontaneous
- ΔG>0 = non-spontaneous.
- ΔG = (ΔH) – T (ΔS)
o Gibbs Free Energy = change in enthalpy – Temperature x change in entropy
- To melt a solid/vaporize a liquid, could add heat so ΔH would be positive.
- For entropy, particles in solids are tightly packed and held together by intermolecular forces
(covalent bonds).
o Liquids have greater degree of disorder.
o Gases are more disordered than liquids.
o Therefore, ΔS increases as we transition from solid  liquid  gas.
- Increase in entropy = thermodynamically favourable, as it contributes to more negative ΔG
- Increase in enthalpy is thermodynamically unfavourable and contributes to more positive ΔG
- Transitioning from solid  liquid  gas is favourable at high temps (maximizes T(ΔS)
o Transitioning from gas  liquid  solid, is favourable at low temps
HEATING CURVES
- Used to depict phase changes as heat is added to substance.
-
-
X-axis = heat added
Y axis = temperature
First slope (blue) is when substance exists as a solid.
o As heat is added, temp of solid increases, and heat goes towards increasing kinetic
energy of particles of solid  vibrate faster, increasing temp.
o Q = MCΔT = how much temperature changes from addition of heat.
 Q = how much heat is added (kilojoules)
 m = mass (kg)
 c = specific heat capacity of a substance (kJ/ kg x K  kilojoules divided by
kilograms x Kelvins)
 NOTE: specific to given substance and to each state.
 ΔT: change in temperature (Kelvins)
Flat  where fusion occurs (where solid melts to liquid) but even though heat is added,
temperature of substance is not changed.
o
-
-
Added heat is being used to break intermolecular forces holding solid together rather
than increasing kinetic energy of the particle.
o Enthalpy of Fusion (ΔHf) = how much heat is needed to melt one mole of substance
(kJ/mole)
After fusion, substance is now liquid and temperature of substance increases again as heat is
continued to be added.
Next part of heating curve is flat, and represents vaporization.
o Heat is added to the substance and disrupts the intermolecular forces between liquid
particles as liquid vaporizes into gas.
o Heat of vaporization (kJ/mole) represents amount of heat required to vaporize one mole
of substance.
Near triangle, adding heat will increase temperature of water vapor.
o If we heat a block of ice until it melts,
HOW MUCH HEAT IS NEEDED TO MELT 20G ICE AT -10 C INTO GAS AT 110 C?
- Any added heat will increase temp until it reaches melting point (0 C)
o Q = MCΔT  know this is right equation b/c delta T = temperature change for -10 and 0
C
o Q = (20g)(2.108 Kj/kg x K)(10 C)  NOTE: make all units consistent
o Q = (0.02kg)(2.108 kJ/kg x K)(10 K)
o Q = 0.42166 KJ of heat
- Next part of heat must go into melting ice into liquid water, which is going to need heat of fusion.
o Q = mΔHf
o NOTE: heat capacity values and enthalpy are provided
o Q = (20g) (334 J/g)
o Q = 6680 J
- Added heat raises temperature of substance which is now liquid water.
o Q = MCΔT
o Q = (20g)4.186 kJ/kg x K)(100C)
o Q = (0.002kg)(4.186 kJ/kg x K)(100K)  from 0 degrees to boiling point of 100 C where
liquid vaporizes into gas
o Q = 8.372 kJ
- Turns into water vapor
o Q = mΔHvap
o Q = (20g)(2256 j/g) = 45120 J
 NOTE: ΔHvap = 2256 J/g
- ADD ALL HEAT VALUES TO OBTAIN TOTAL HEAT NEEDED
o Q = 0.4216 kJ + 6680 J + 8.372 kJ + 45120 J + 0.3992 KJ
o Q = 421.6 J + 6880 J + 8372 J + 45120 J
o Q = 60 992.8 J or 61KJ
PHASE DIAGRAMS
- phase diagram tells us state of matter when
substance is the most stable at given temperature
and pressure.
- Diagram shows phase diagram for water.
- Left (light blue) represents solid phase.
- Bottom (yellow) represents gas phase.
- Right (teal) represents liquid phase.
- Boundary lines separate 3 phases
- At boundary lines, 2 phases are in
-
-
equilibrium
- Fusion and freezing occur when solid liquid
boundary line is crossed during transition
from solid to liquid (vice versa)
- Vaporization
and condensation are
equivalent concepts for liquid to gas boundary line
Sublimation and deposition are equivalent for solid gas transition.
ONLY FOR WATER = solid liquid boundary slants up into left while for other substances,
would slant into the right
o Due to water being less dense in solid than liquid form due to complex crystal
lattice structure
o At low temps, increasing pressure on ice will cause it to melt.
For CO2, liquid carbon dioxide doesn’t exist at atmospheric pressure so solid CO2 (dry ice)
sublimates readily into gastro atmosopheric pressure.
SPECIAL POINTS
Triple Point: represents temperature and pressure
at which all 3 phases are at equilibrium and substance is
combination of all 3 phases.
Critical Point: terminal point of liquid gas boundary line,
-
which represents temperature and pressure conditions beyond
which liquid and gas phases become indistinguishable.
o
Beyond critical point, fluid is no longer a fluid or gas
but a supercritical fluid.
Supercritical fluid have low viscosity of a gas but
high desnity of a liquid, giving them properite sof both.
STOICHIOMETRY AND BASIC CHEMISTRY MATH
CHEMICAL FORMULA AND PERCENT COMPOSITION
CHEMICAL FORMULAS
- Molecular: the type and number of each atom in a molecule. Ex. C41010, H2O, C6H12O6
- Empirical: mathematically simplified version of molecular formulas; states only the ratio of atoms
in a compound. Ex.C2H5, C6H12O6  CH2O
o Ex. hydrocarbons with 1 double bond (propene, octene) have formulas that follow
patterns of CxH2X
o Propene has molecular formula of C3H6 and octene has molecular formula of C8H16 but
both of their empirical formulas are the same – CH2
- Empirical helps us identify molecules that share common atomic ratios, but share properties that
make them fall into the same categories.
o Limitation: tell us nothing about total number of atoms in molecule or overall weight.
EXAMPLE
What’s the empirical formula for a 150g sample composed of 60g of carbon, 10g of hydrogen, and 80g of
oxygen?
- Convert samples to moles by dividing by atomic weight.
- C = 60 x 1 mol / 12 = 5 mol
- H = 10 x 1 mol / 10 = 10 mol.
- O = 80 x 1 mol / 16 =5 mol.
Given the molecular formula CH2O, what is the molecular formula if the molar mass of the sample is 180
g/mol?
- To find molecular formula, need to have molar mass.
- C: 1 x 12g/1 mol = 12 g/mol C
- H: 2 x 1 g/1mol = 2 g / mol H
- O 1 x 16 g/1 mol = 16 g.mol O
- Total mass of CH2O = 20g.
- 180 / 30 = 6, need to multiply the molar ratio by 6 to obtaib molecular formula
- Molecular formula = C6H12O6
PERCENT MASS
What is the amount (in grams) of each individual element of 150g compound that contains 40% carbon,
7% hydrogen, 53% oxygen?
- Carbon: 40% = 150 x 0.4 = 60g
- Hydrogen: 7% = 150 x 0.07 = 10g
- Oxygen: 53% = 150 x 0.53 = 80g
- Then need to convert if there was 100g of the sample
o C: 40% = 100 x 0.04 = 40
o H: 7% = 100 x 0.07 = 7g
o O: 53% = 100 x 0.53 = 53
- Then, convert to mols.
o C: 40 x 1 mol/12g = 3.3
o H: 7 x 1 mol/1g = 7
o O: 53.x 1 mol/16g = 3.3
PERCENT COMPOSITION BY MASS
What is the percent composition by mass of oxygen in a molecule with empirical formula CH2O?
- C: 1 x 12g/1mol = 12 g/mol
- H: 2 x 1g/1mol = 2g/mol
- O: 1 x 16g/1mol = 16g/mol
- Molar mass of CH2O = 12+2+16 = 30g/mol
-
C: 16/30 x 100 = ~50
H: 1 / 30 x 100 = 3.33%
O:
CHARACTERISTICS OF GASES AND IDEAL GASES
- Ex. CO2 and O2 are in respiratory and circulatory system
- NO as cell signalling molecules.
- Gases are compressible whereas liquids and solids are incompressible (will always have the same
volume)
- Gases can expand to fill a given space or be compressed down into a smaller container.
KINETIC MOLECULAR THEORY
- System of tiny particles bouncing around inside given region of space.
- Asserts that pressure of walls on container is result of elastic collisions of particles within those
walls.
- Average kinetic energy of particle is directly proportional to the temperature of the gas.
- Brownian Motion: random particle movement, which was due to result of random bouncing
around the molecules in medium.
- Ideal Gas:
o assumes that gas particles don’t have any volume (no size) (ex. if container has 5L, it is
the space between the molecules because the molecules themselves don’t take up any
space.
o Gas particles experience no attractive or repulsive force. (no hydrogen bonding, no
dipole-dipole)  if they pass by eachother, no interaction unless they elastically bounce
off eachother.
o Collision are perfectly elastic, motion is random, time in contract during collision is legible.
IDEAL GAS BEHAVIOUR
Favour ideal behaviour that are generally more favourable to gaseous states.
- Ex. liquids like to boil to gas when they’re hot, so  temperature, more likely to be ideal gas, 
volume with big space with  pressure.
- As temperature increases, intermolecular forces can be overcome and thermal energy increases.
PRESSURE, VOLUME, TEMPERATURE
BOYLE’S LAW
Pressure of volume of gas are inversely related.
Can express inverse relationship in 2 ways:
P1V1 = P2V2  P1,V1 are pressure and volume of gas in 1, P2,V2 pressure
and volume of gas in another.
CHARLES’S LAW
Volume and temperature of gas are directly related.
Ex. can be a problem, if you are in hot area and have tires overinflated in the heat, gas inside expands and damage/explode the tire.
AVOGARDO’S LAW
- Volume of gas is directly related to number of moles of gas particle.
- v = volume
- N= moles
- 1 and 2 = 2 states of same system.
- Can also be written as V/N = K
- In any 2 containers of gas that are the same size, will have same number of gas
particles, assuming they are kept in same temperature and pressure.
- At 273K and 1 atm, that mole of gas take sup 22.4 L of space, called molar volume.
o Standard temperature and pressure (STP): pressure = 1atm, temperature = 0 C or 273K,
conditions commonly used for calculations involving gases.
- Standard Thermodynamic Conditions: 1 atm and 25 C (298 K)
IDEAL GAS LAW AND REAL VS IDEAL GASES
- IDEAL GAS LAW: PV = nRT
o R = ideal gas constant  0.08: atm/mol K or 8.31 J/mol K
- Testing the Ideal Gas Law
Plug and Chug questions
Setting up the Equation
Relationships
If pressure is doubled and volume is halved, what happens to temperature of gas?
- Need to plug the changes into ideal gas laws
- New equation = 2P(1/2V) = nRT
REAL GASES
- Need to acknowledge that gas particles have a size and there might be some intermolecular forces
between them.
- Need to use Van Der Waals equation.
- Vm = molar volume
- A = constant for each gas
- A = attractive intermolecular forces
- B = size of particle
- Really polar molecule = large a term and will behave nonideally.
EX. 10 L container with 0.5 mols of gas, would real pressure inside container be smaller or larger than the
ideal pressure?
- Answer depends on if we are talking about normal(ish) conditions or extreme conditions
-
-
-
Under normal conditions such as STP: real gases have a slightly smaller but nearly equivalent
pressure when compared to an ideal gas.
Attractive forces of particles can act across long distances.
Even when gas molecules are further away from eachother (dipole dipole interactions can cause
them to clump and stick to eachother rather than pressing out on container.
Can result in a lowering P.
Under extreme pressure or low temperatures: real gases have higher pressure when compared
to an ideal gas at extreme temperatures.
o Higher pressure = more volume, resulting in higher P.
Real gases have a lower pressure when compared to an ideal gas at low temperatures.
IDEAL VS REAL GASES
- Ideal gas law assumes that molecule has no size and no attraction towards eachother.
o PV = nRT
- Real gases use Van der Waals to account for attractive forces between the molecules and how big
the molecules are
o (P+ a/V2m)(Vm-b) = RT
- Normal conditions: real gases have smaller but nearly equivalent pressure when compared to an
ideal gas.
- Extreme conditions: real gases have higher pressure when compared to an ideal gas at extreme
temperatures.
- Extreme conditions: real gases have lower pressure when compared to an ideal gas at low
temperature.
NOTE: when there is increased pressure (and all other variables are held constant), temperature and
number of moles are increased and volume is decreased.
DALTON’S LAW AND GRAHAM’S LAW
DALTON’S LAWS OF PARTIAL PRESSURE
- Total pressure of mixtures of gases is equal to the sum of partial pressures of individual gases.
- Pressure that gas exerts is independent of any other gases present and depends only on total
pressure and mole fraction of gas.
- Ptotal = Pa +Pb+Pc…
- Mole Fraction eq: Xgas = ngas/ntotal
- Ideal Gas Law rearranged: n = PV/RT
-
Mole fraction final: xgas = Pgas/Ptotal
To get Dalton’s law, need to rearrange expression to get pressure of gas (partial pressure) is:
Pgas = (xgas)(Ptotal)
o Note: Dalton’s law assumes the particles aren’t interacting with eachother and are of
negligable size.
GRAHAM’S LAW OF EFFUSION
-
-
When a gas is effusing (escaping form small opening) smaller, lighter particles escape faster and
the big, massive particles move slower, escaping at a slower rate.
Square root comes b/c movement of particles can be described with kinetic energy equation:
KE= ½ mv2
Since gases are all mixed together, they are at the same temperature. (temperature is
proportional to average kinetic energy)
o Average Kinetic Energy (U) = 3/2kT
Proportion of rate of movement is inversely related to square root of proportion for molecular
mass.
THERMODYNAMICS AND GIBBS FREE ENERGY
CALORIMETRY
- Avg candy cane = 50 cals
- Calorimeter: measures heat energy produces by burning that substance.
o Mimics combustion – how body processes nutritional calories.
- Nutrition molecules in food undergo combustion rxn in body – molecules consumed by presence
of O2, producing H2O and CO2 – body gets energy from this.
- Calorimetry – substance undergoes rxn in well insulated container and temp change measured to
estimate heat released during rxn using bomb calorimeter.
UNITS OF HEAT
- Joules (J)
- Calories (cal) = 4.184 J
o 1 cal = amount of energy required to raise temp of 1 gram of water by 1 degree C
o Can use to see how much heat was produced with Q = mcΔT
- Nutritional calorie (Cal)
o 1 Cal = 1000 cal= = 1kcal
- Q = mcΔT
o m= mass of water involved
o ΔT= change in temperature
o c= specific heat capacity (of water)  1mL H2O = 1g
- Specific Heat Capacity (c): energy required to raise temperature of 1 gram of substance by 1
degree C.
o
Water heat capacity = 4.186J per g x C
ENTHALPY
- Enthalpy (H): total heat in thermodynamic system – equal to internal energy of given system +
product of its pressure and volume.
- Expressed in joules
- ΔH = ΔU + PΔV
o H = enthalpy
o U = internal state of system
o P = pressure
o V= volume
- PΔV – pressure volume work is form of work associated with expansion/compression of system.
- ΔU – internal energy = heat put into system (Q) and work done on system (or minus work down
in surrounds)
o ΔU = Q – work(surroundings)
- Q = ΔU + W  heat added to system = internal energy change + work done by system.
EXOTHERMIC VS ENDOTHERMIC
- Endothermic Process: If heat added to system, system gains heat, ΔH is +, ΔH>0
o Require heat input.
- Exothermic Process: heat is removed from a system and system loses or produces heat, ΔH is (-),
ΔH<0
o Release heat to environment.
- Ex. chemical rxn occurs btwn 2 molecules in aqueous solution lowers temp of surrounding water.
Endo or exo?
o Did the system (molecules participating in rxn) gain or lose heat?
o The environment lost heat and got colder but this means that chemical reactants must
have sucked up heat from environment to obtain heat needed for reaction to occur.
o Endothermic process  heat taken up by rxn can be used to break bonds, enable
chemical rxn to take place.
- Ex. if we take an ice pack form the freezer and it feels cold. Is it endo or exo?
o Chemicals in ice pack stealing heat from their environment so it’s endothermic.
- Ex. is fire endo or exo?
o Fire is result of combustion rxn that releases large amount of heat into environment –
Fuel + O2  H2O + CO2
o If you feel heat, the system is losing heat so it’s exothermic.
STANDARD ENTHALPY
- ΔH = enthalpy change under standard conditions (1atm, 25C)
o 25 degrees b/c most rxn performed near room temperature.
- Standard Enthalpy of Formation/Standard Heat of Formation (ΔHf) – enthalpy change associated
with forming one mole of compound from its elements under standard conditions.
- Ex. CO2 can be formed from carbon and oxygen gas
o When combine, to form one mole of CO2, this rxn releases 394 kJ of energy (exotermic
rxn)
- Standard enthalpy of any element in its standard state if 0
STANDARD ENTHALPY OF REACTION
-
ΔH rxn = ΔH products - ΔH reactants
Ex. combustion of methane
- Need to know standard enthalpies of formation of each reactants and products.
- CH4 + 2 O2  CO2 + 2H2O (values provided)
BOND DISASSOCIATION ENERGY
Enthalpy Change Associated with Breaking Bonds
- Formation of bonds = exothermic
- Bond breakage = endothermic
- Δ H rxn > 0 = endothermic
- Δ H rxn < 0 = exothermic
- Average Bond Enthalpies: enthalpy charge associated with breaking of each of these bonds.
ENTHALPY THROUGH MULTI-STEP RXN
- Hess’s Law: total enthalpy change for overall process is the sum of enthalpy changes in each step
RULES FOR ADDING ENTHALPY RXN
- Each step must be ordered in proper forward direction with respect to full rxn.
- Reverse sign of ΔH degrees when reversing order of component step.
- When changing stoichiometry, multiply or divide enthalpy change by same coefficient.
- Need to reverse sign when reversing order of component step.
ENTROPY AND GIBBS FREE ENERGY
- SPONTAENEITY OF RXN: whether a rxn will occur
- Spontaneous rxn is one that will proceed on its own
- Non Spontaneous Rxn is when the reaction won’t proceed without an outside energy source.
GIBBS FREE NEERGY (ΔG)
- Energy that can be used to perform work in reversible rxn
for reaction to take place, activation energy (Ea) must be
overcome. (above red line)
Change in ΔG as a whole only takes into account the starting state
of reactants and final state of product.
If free energy decreases over rxn (products in more stable state
than reactants), then rxn is wanting to occur and be spontaneous.
Decrease in ΔG (or negative ΔG change) is associated with
spontaneous reactions, also called exergonic rxn ΔG<0
- X tells us that ΔG is lost or decreases
- if products are higher in energy than the reactants, rxn less likely
to process, ΔG+, nonspontaneous, endergonic rxn.
FREE ENERGY STATE OF SYSTEM
Determined by:
- Enthalpy – heat transfer between systems and surroundings where exothermic rxn release heat
and (-) and endothermic rxn absorb heat (+)
- Entropy – measure of disorder, tells us about number of configurations system can have (greater
number = higher entropy)
o Gas has higher entropy than liquids do (solids most constrained)
o Larger molecule breaking to smaller multiple molecules produce entropy increase and so
does unfolding of proteins.
o Predict how likely chem rxn is perceived spontaneously
- Temperature
- Spontaneity is associated with -ΔG and with exothermic rxns
o ΔG = ΔH – TΔS
Ex. ΔH is 14 kJ and ΔS = 100 J/K, under what conditions
will it be spontaneous?
Since signs both positive, can plug values into Gibbs
Free Energy eq.
-
NOTE: convert units (ΔH 14 kJ to 14 000 J)
Solve for temperature, set ΔG to 0
Ans: 140 K (higher than 140, TΔS more + and ΔG –
Rxn is spontaneous at temperatures greater than 140K
STANDARD GIBBS FREE ENERGY CHANGE
- ΔG: Gibbs Free Energy change under standard condition (1 atm, 25 C, 1M concentration)
- ΔGf = Gibbs Free Energy of formation
o ΔGf = 0 (in standard state)
- ΔGrxn =  ΔGproducts -  ΔGreactants
THERMODYNAMIC AND KINETIC CONTROL OF RXN
- Equilibrium  endpoint of reversible chemical reaction – rate (forward) = rate (reverse)
THERMODYNAMICS VS KINETICS
- Thermodynamics – where rxn wants to be. Lots to do w/ stability.
o Concerned with stability, spontaneity (ΔG) and equilibrium.
- Kinetics – how fast the reaction moves in that direction.
o Concerned with reaction rate, activation energy, reaction mechanisms.
o Activation Barrier: amount of activation energy that must be overcome for rxn to go to
completion
 Barrier is high b/c it represents transition state: temporary configuration that’s
unstable and high energy
 State that reactant must transition through to become product molecules.
 Reactant  transition state  product.
EQUILIBRIUM CONSTANT
- ΔG and Keq (equilibrium constant)
o Keq = [products]/[reactants]  thermodynamica concept
o ΔG = -RT In Keq
 -R = gas constant
 T = temperature
 In Keq = natural log of Keq
 eq useful as it relates spontaneity of rxn under Stannard conditions
- if Keq > 1, products favoured over reactants, ΔG will be negative and rxn will be spontaneous in
forward direction from reactants to products (products favoured at equilibrium) [reactants 
products]
- If Keq < 1, reactants favoured over the product, ΔG+ and rxn is non spontenaous in forward
direction, want to proceed in reverse direction (reactants  products) [reactants favoured at
equilibrium)
-
If Keq = 1, natural log of one is 0 and Gibbs free energy change will be equal to 0. Forward and
reverse reaction are equally favourable from thermodynamic perspective (reactants  products)
At equilibrium, Rate (forward) = Rate (reverse)
o k(1) reactants = k-1 (products)
o k1/k-1 = [products][reactants]
NONSTANDARD CONDITIONS
- ΔG = ΔG + RT In Q
-
-
o
Q = reaction quotient may be sued for reactant and product concentration in any point in
rxn.
When Q > Keq,
so reaction would want to proceed in reverse to get
equilibrium,
When Q < Keq, product to reactants lower than it would like to be at equilibrium so rxn wants to
proceed in forward direction from reactants (reactants  products)
Q = Keq, we are at equilibrium and ΔG=0
CATALYST
- Catalyst increases the rate of rxn by lowering activation energy.
- Increases rate of rxn in forward and reverse directions
- Doesn’t affect spontaneity.
- Catalyst affect how fast rxn reaches equilibrium
- Affect kinetics but not thermodynamic (spontaneity, stability, equilibrium state of rxn)
ACTIVATION ENERGY AND REACTION COORDINATES
CHEMICAL KINETICS
- For rxn to occur, reactants must collide with eachother in correct spatial orientation and with
sufficient kinetic energy for rxn to take place.
When energy threshold is reached, reactant form short lived and highly
unstable transition state or transition complex before stabilizing into ultimate
products of rxn.
Chemist visualize kinetic energy using reaction coordinate diagram
which displays the energy profile of rxn as it progresses.
o
X axis = time (reaction progress), while Y-axis = Gibbs Free Energy
Spontaneous rxn are associated with -ΔG and is spontaneous.
-
If products have greater energy (higher on diagram than reactants), rxn has +ΔG value and is nonspontaneous.
ACTIVATION ENERGY
- Energy threshold that must be reached for rxn to proceed.
- On rxn coordinate, Ea is the difference in energy between initial reactants
and peak of the graph, representing the energy input required to reach
transition state complex.
- Reactants must acquire enough energy to surpass energy barrier before
rxn can proceed to form products
- Higher Ea = harder for rxn to occur = slower rxn
- Arrhenius Equation:
higher activation energy corresponds to lower rxn rate.
Ea should be inversely proportional to rate constant.
The higher the Ea, the more negative the number and
rate will be
Temperature directly proportional to rate constant (high
temp = higher kinetic energy over Ea barrier)
-
2 ways that reaction rate can be increased
o Lower activation energy
o Increased temperature
ENZYMES
- Biological catalyst reduce Ea of a reaction.
o Stabilizing transition state
o Weakening bonds within the reactants
o Changing orientation of reactants to facilitate effective collision
o Increasing frequency of collisions
o Donating electron density to reactants
- Increase the rxn rate
- Do not affect thermodynamic properties of rxn (ΔG, enthalpy, entropy) and cannot turn nonspontaneous rxn into spontaneous
o Not consumed in rxn
o Little amount of catalyst  large amount of product
HETEROGENEOUS AND HOMOGENOUS CATALYSIS
- Based on phase of catalyst compared to phase of reactant species
- Heterogeneous catalyst: different phase than reactants
- Homogenous catalyst: same phase as reactant