1.1 Atoms, Molecules and Chemical Formulas
1.2 Measurements in Chemistry
Physical Properties: size, colour and mass
Quantitative Measurements of properties
Every Measurement gives a numerical result of 3 aspects:
Precision: exactness of a measurement
Accuracy: close a measurement is to the true value
Chemical Properties: chemical transformations
1.3 Chemical Problem Solving
1.4 Counting Atoms: The Mole
Mole and Avogadro’s Number
Avogadro’s Number : 6.022 X10 23 items/mol
Molar mass: (Fractional Abundance) (isotopic molar mass)
Mass- Mole-Atom Conversions:
1.5 Amounts of Compounds
1.6 Aqueous Solutions
Solvent: substances used to dissolve solutes
Solute: pure substance dissolved in solution
Polyatomic ions remain intact
Concentration (c ) amount of solute in a solution
Any solutions is formed by dissolving an ionic solid in water conducts electricity
Dilutions: amount of solute remains the same but the volume of the solution
RESULTS IN A : solution of lower molarity
The number of moles of solute does not change during the dilution
C1 1= C2 2 Concentrated acid is always added to the water
1.7 Writing Chemical Equations
The number of atoms of each element is conserved in any chemical reaction
1.8 The stoichiometry of Chemical Reactions
Stoichiometry is the study of the amounts of materials consumed and produced in
1.9 Yields of Chemical Reactions
Under practical conditions, chemicals almost always produce smaller amounts of
products than the amount predicted by stoichiometric analysis.
1. Many reactions stop before reaching completion ( do not go to completion
because they reach dynamic equilibrium)
2. Competing reactions often consume some of the starting materials
3. When product of a reaction is purified and isolated, some of it is inevitably
lost during the collection process (impossible to get all the product out of the
1.10 The Limiting Reactant
The reactant that runs our is called limiting reactant ( imits the amount of product
that can be made). Other starting materials are the excess reactant
Limiting reactant: smallest value of moles divided by coefficient
The pressure of the atmosphere can be measured with the instrument called the
Units of pressure : standard atmosphere ( atm) pressure that will support a column
of mercury 760mm in height.
Manometer: gases exert pressure on both liquid surfaces. The difference in pressure
exerted by two gases.
2.2 Describing Gases
Volumr of the trapped gas is inversely proportional to the total pressure applied bu
mercury plus the atmosphere.
V 1/Pgas ( fixed temperature and amount) Volume of a gas is directly proportional to its temperature:
Vgas Tgas ( fixed pressure and amount)
Gas volume is proportional to the amount of gas
Vgas ngas (fixed pressure and temperature)
Ideal Gas Equation
Universal Gas constan: R= 8.314 LkPa/mol K or 0.08314 L bar/mol K
Ideal Gas law: pV=nRT
As pressure increases= Volume increases
Solving quantitative problems about gases at moderate temperatures and pressures
requires only one equation. The ideal gas equation.
2.3 Gas Mixture
Dalton’s law of partial pressure
Ptotal 1p 2p +3 4
(each p is an individual pressure)
Mole Fraction (X):
Parts/million (ppm) and Parts/billion (ppb)
PA= XA Total
In a mixture of gases, each gas contributes to the total pressure the pressure that it
would exert if the gas were present in the container by itself.
2.4 Gas Stiochiometry
Liquid or solid n= m/M
Aqueous Solution n=cV
2.5 Molecular View of Gases
E = 3RT/2N
At a given Temperature, all gases have the same kinetic energy distribution
Ideal gas: negligible molecular sizes
Negligible intermolecular forces The volume occupied by the molecules of an ideal gas is negligible compared with
the volume of its container. The energies generated by forces among ideal gas
molecules are negligible compared with molecular kinetic energies.
2.6 Additional Gas Properties
Root Mean square :
Diffusion: movement of one type of molecule through molecules of another type
Effusion: movement of molecules escaping from a container into a vacuum
2.7 Non-ideal (real) Gases
Van der Waals Equation:
2.8 Chemistry of the Earths Atmosphere
NO 2 red-brown gas that can be seen in the atmosphere over large cities, absorbes
energy from sunlight and decomposes into NO molecules and oxygen molecules)
O 3 SO 2 SO 3
Photochemical Smog: mixture of all these pollutants
Dynamic Equilibrium:rate of evaporation equals the rate of condensation
Vapour Pressure of any substance increases rapidly with temperature because the
kinetic energies of the molecules increase as the temperature rises.
Humidity= (100%) (pH2o/pvapour,H2O)
3.1 Types of Energy
Energy: ability to do work
Work: displacement of an object against an opposing force
Kinetic Energy : moving objects
Potential Energy: Stored energy
Thermal Energy: Content of a hot object
Radiant energy: content of electromagnetic radiation (light/ infrared radiation)
Kinetic and Potential Energies
SI unit: kgm /s Potential energy= gravitational energy
Electrical forces lead to electrical potential energy. Energy is released when
oppositely charged objects move close together whereas energy must be supplied to
pull oppositely charged objects apart.
Total energy of these random movements
Average molecular kinetic energy increases as temperature increases.
Liquid phase atoms and molecules still have energies of translation, rotation and
Thermal energy is the cause and temperature is the effect.
Energy transfers and transformations:
When atoms/molecules collide with one another, energy transfers cause some
molecules to speed up and others to slow down.
Energy transformations accompany chemical reactions:
Reaction releases energy, chemical energy is converted into other forms of energy.
- Quantitative study of energy transfers and transformations
System: describe and study by it self
Surroundings: everything else
System is separated from its surroundings by a boundary across from which
matter/molecules can move
Conservation of Energy
Law: energy is neither created nor destroyed
A system can exchange thermal energy with its surroundings. Amount of energy that
is transferred is heat (q). Measured in Joules.
∆T depends on the q the amount of heat transferred, direction of the heat flow,
∆T absorbs heat :+ ∆T releases: -
∆T depends on the identity of the material. Expressed by Molar Heat Capacity
q surr-q sys Work
Energy used to move an object against an opposing force (w)
Amount of work depends on the magnitude of the force that must be overcome and
the amount of movement or displacement. w=Fd
W surr-W sys
First Law of Thermodynamics
Energy(E) is exchanged in two ways : heat or work
∆E sysE surr
Law: ∆Etotal=∆Esys +∆Esurr=0
State and Path Functions
only on the conditions that describe the system before and after the transformation
State Function: How the changes occur
Path Function: how a change takes place
Key concept: the value of a state function does not depend on the path taken or on
the rate of the change.
Thermodynamic Path Function
Energy is a state function
Heat and Work are Path Functions
3.3 Energy Changes in Chemical Reactions
Energy is released during the reaction
Reversing the direction of the reaction changes the sign of the energy changes.
Features of Reaction Energies
When a chemical reaction releases energy. ∆E has a negatives sign.
Reaction absorbs energy. ∆E has a positive sign.
∆Ereaction= ΣBEbonds broken – ΣBEbondsf formed
3.4 Measuring Energy Changes: Calorimetry
Calorimeters : measure heat flow that accompany chemical processes.
Chemical reaction takes place within a calorimeter, resulting in a heat flow between
the chemicals and the calorimeter. The temperature of the calorimeter rises or falls
in response to this heat flow.
qchemicals = -qcalorimeter
release heat, qchemicals is negative EXOTHERMIC
gains heat, qcalorimeter is positive , temperature rises
Absorbs heat, qchemicals is positive ENDOTHERMIC
Loses heat(qcalorimeter is negative) , temperature falls qcalorimeter =Ccal∆T
Calculating Energy Change
Enthalpy (H)= Heat flow at constant pressure
Hess Law: Overall ∆H= Sum of ∆H for individual steps
Reactants (net reaction) Products
(decomposition reactions) Elements (formation reactions) Products
Standard Enthalpy of formation (∆H )f formation reaction
Formation reaction= 1 mol of a substance produced from elements
1/2N 2g) +1/2 O 2g) N(g)
∆H of reactioneff products∆H fproductscoeff reactants ∆Hof reactants
∆Hreaction= ∆E reaction + RT∆ngases
A formation reaction produces 1 mole of a chemical substance from the elements in
their standard states.
The enthalpy change for any overall process is equal to the sum if the enthalpy
changes for any set of steps that leads from the starting materials to be products,
Enthalpy is a thermodynamic state function that describes heat flow at constant
4.1 Characteristics of Atoms
Fundamental characteristics of atoms:
Atoms possess mass:
Atoms contain positive nuclei
Atoms contain electrons
Atoms occupy volume
The volume of an atom is determined by the size of its electron cloud. Atoms have various properties
Atoms attract one another
Atoms can combine with one another: form chemical bonds with one another to
4.2 Characteristics of light
Studying the structure of atoms by electromagnetic radiation. Light is a form of this
Light has wave aspects
Light is a wave like property.
Frequency(v): number of wave crests passing a point on space in one second.
Unit s or Hz
Wavelength (λ): distance between successive wave crests.
Units m or nm
Amplitude: height of a wave
Amplitude of a light wave measures the intensity of the light.
Total energy of a beam of light depends on its intensity.
Photoelectric effect : energy of light depends on its frequency and intensity.
Light comes on bundles of photons. Eeach photon has an energy that is directly
proportional to the frequency.
Ephoton: energy of light
E=hv , v=c/λ , E= hc/λ
Electron kinetic energy= Photon energy –Binding energy.
Higher intensity means more photons but not more energy per photon
Light has particle Aspects
Light has some properties of waves and some properties of particles.
4.3 Absorption and Emission Spectra
when light interacts with free atoms , interactions reveal infomormation about
electrons bound to individual atoms.
Light and energy
Absorption of photons by free atoms has two different results, depending on the
energy of the photons.
When an atom absorbs a photons suffieciently high energy, an electron is ejected,a
process called photoionization.
2 types of results:
Transferred to a higher energy state Excited State
-give up their excess energy to return to loswer energy states lowest energy state of an atom, which is the most stable state groung state
Attractive electrical forces hold a bound of electrons within an atom ad energy must
be supplied to remove a bound electron from an atom.
Lower the energy state, more energy must be supplied to remove the electron.->
energy changes that are measured relative to the energy of free electrons. Energy of
a free stationary electron is zero.
Exchange of energy between atoms and light is that energy is conserved. Therefore,
change in energy of the atom exactly equals the energy of the photon
When an atom absorbs a photon, the atom gains the photons energy so ∆Eatom is +
Atom emits a photon, atom loses the photons energy so ∆Eatom is –
As an atom returns to its ground state it must lose exactly the amount of energy that
is originally gained,However the excited atoms lose excess energy involving small
energy changes do the frequency of emitted photons are often lower than absorbed
Absorption spectrum: unique for each gas because different atoms absorb different
energies of photons.
Emission spectrum: intensity of light emitted as a function of frequency.
Each Frequency absorbed or emitted by an atom corresponds to a particular energy
change for the atom.
Quantization of Energy
When an atom absorbs light of frequency v, the light beam loses energy hv, and the
atom gains the amount of energy.
Therefore, a photon with high enough energy can cause an atom to lose one of
electrons.Absorption of a photon recults in an energy gain for an electron in the
atom energy change for the atom equals the energy change for an atomic electron.
En= 2.18x10 J/ n 2
4.4 Properties of Electrons
Electrons have particle like and wave like properties.
-Each electron has the same mass and charge.
-Electrons behave like magnets: occurs due to property called spin.
-Electrons have wave properties.
The momentum of a particle is the porfuct of its mass and speed; p=mv
Λparticle= h/mν 4.5 Quantization and Quantum Numbers
Each atomic energy level is associated with a specific three-dimensional atomic
Principal Quantum Number (n):
- indexes energy.
Each electron in an atom can be assigned a value n that is a positive integer that
correlates with the energy of the electron.
Most stable energy for an atomic electron corresponds to n=1 and each one higher
value of n is less stable energy state.
-Size of an atomic orbital: energy of an electron correlated with its distribution in
-Higher the principal quantum number the more energy the electron has and the
greater its average distance from the nucleus.
n: Increases, energy of the electron increases, orbital gets bigger, electron is less
tightly bound to the atom.
Azimuthal Quantum Number (l)
-Shape of atomic orbitals
-l correlates with the number of preferred axis in a particular orbital and thereby
identifies the orbital shape.
Value of l 0 1 2 3
Orbital letter s p d f
Magnetic and Spin Orientation Quantum Numbers(m) (m ) l s
-magnetic quantum number m: indeles restrictions
l=1 m:l-1, 0, +1
-electron has magnetism associated with a property called spin.
-Magnetism is directional, so the spin of an electron is directional
-Spin orientation is quantisized: electron spin in 2 ways ; UP or DOWN
-Spin orientation quantum number m : indsxes the behavior
-m sre +1/2 and -1/2
4.6 Shapes of Atomic Orbitals
-An atom that contains many electrons can be described by superimposing(adding
together) the orbitals for all of its electrons to obtain the overall size and shape of
-Chemical properties of an atom are determined by behavior of their electrons
because atomic electrons are described by orbitals, the interactions of electrons can
be described in terms of orbital interactions
-Two characteristics of orbitals determine how electrons interact : SHAPES &
-Quantum number n&l determine the size and shape of an orbital.
-n increases the size of the orbital increases
-l increases, shape of the orbital becomes more elaborate Orbital Despictions
3 types of plots :
Electron density plot : represents the electron distribution in an orbital as a two
dimentional plot. Does not show it three dimentional
Orbital Density picture: Take a lot of time and care
Electron contour drawing: all details of electron density inside the surface is lost
-Value for r where electron density drops to zero called Node
-Electron has wave like properties- this wave has zero amplitude at the node, but
non-zero amplitude on either side of the node. Thus, electrons can move across the
node without actually existing at the node.
-Number of nodes increases as s and n increases. However 1s orbitals have no
nodes, a 2s orbital has one node.
In any particular atoms;
-orbitals get larger as the value of n increases
-all orbitals with the same principal quantum number are similar in size
-Each orbital becomes smaller as nuclear charge increases; As the positive charge of
the nucleus increases, the electrical force exerted by the nucleus on the negatively
charged electrons increase as well, and electrons become more tightly bound. Thus
reducing the radius of the orbital, orbitals shrink in size as atomic number increases.
4.7 Sunlight and the Earth
Thermosphere (heat released) -85km molecules of nitrogen and oxygen absorb x-
ray radiation coming from the sun.
N2+hvN2 +e-+ O2+hv O+O
Reactions are unstable and will recombine releasing energy in the form of heat;
N2 +e-N2 +heat O+OO2 +heat
-Aurora : atoms and molecules that reach excited states by gaining energy from
some external source return to their ground states by emitting photons.Absorption
by atoms and molecules removes high energy photons.
-Result: intensity of high energy light decreases as sunlight moves down through the
atmosphere towards the Earth’s surface
Reactions in the Ozone Layer
Stratosphere, solar radiation generates an abundance of ozone (O3)
-Ozone layer forms in two steps; First photon with a wavelength between 180-
240nm break on O2 molecules into two atoms of oxygen.
-Second, oxygen molecules capture one of the oxygen atoms to form an ozone
O2+O O3 +heat
-Second step occurs twice for each O2, giving the overall balanced process for ozone
3O2+hv 2O3 +Heat Why id the ozone layer confined to one region of the atmosphere?
-production of ozone requires both a source of oxygen atoms and frequent collsions
between the atoms and the molecules that make up the atmosphere.
Lower than 20 km O2 dissociation does not occur
Higher than 35km plenty of light to dissociate O2, but molecular density and rate od
molecular collisions are too low.
Interactions of molecules and light in the ozone layer result in a delicate balance
that holds the ozone concentration at a relatively constant value;
-photons with wavelengths 180-240nm break apart O2molecules
-photons with wavelengths of 200-340nm break apart O3 molecules
-oxygen atoms combine with O2 molecules can produce O3 molecules and heat
Gases in troposphere: carbon dioxide, water vapor and methane
Known as Green house gases.
-Glass panes of greenhouses heep the greenhouse warm, greenhouse gases
moderate temperature changes from night to day by absorbing some of the infrared
photons emitted by the earth. Following the absorption the gases reemit still longer
wavelengths photons some where plants reabsorb them.
-Out coming: infrared radiation emitted by the Earth
5.1 Orbitals Energies
A hydrogen atom can absorb a photon and change from its most stable state (ground
state) to a less stable state (exited state).
The Effect of Nuclear Charge
The stability of an orbital can be determined by measuring the amount of energy
required to remove an electron completely:
Ionization energy (IE):
HH +e- IEH: 2.18x10^-18 J
He He +e- IE : 8.72 x10^ -18 J
LARGER THE IE HAS STRONGER GROUND STATE AND STABLE.
Effects of Other Electrons
Multi electron atom, each electron affects the properties of all the other electrons. A
given orbital is less stable in a multielectron atom than it is in the single electron ion
with the same nuclear charge.
This electron- electron repulsion cancels a portion of attraction between the nucleus
and the incoming electron. Partial cancellation is called screening. -incomplete screening can be seen in the ionization energies of hydrogen atoms,
helium atoms and helium ions.
-Without screening , the IE of a helium atom would be the same as that of a helium
-With complete screening, one helium electron would compensate for one of the
protons in the nucleus
-Electrons in compact orbitals pack around the nucleus more tightly than do
electrons in large, diffuse orbitals.
-RESULT: the effectiveness in screening the nuclear charge decreases as the orbital
-Size of an orbital increases with n, and an electron ability to screen decrease as n
-Higher the value of the quantum number l, the more that orbital is screened by
electrons in smaller, more stable orbitals
- Electrons with the same l value but different values of l do not screen one
another effectively .
-ex. Electrons occupy different p orbitals that have the same n value = same energies
known as degenerate, the amount of mutual screening is slight
-Quantitative information about energies of atomic orbitals is obtained using
photoelectron spectroscopy principles of the photoelectric effect to gaseous
5.2 Structure of the Periodic Table
Pauli Exclusion Principle: EACH ELECTRON IN AN ATOM HAS A UNIQUE SET OF
AUFBAU Principle: ELECTRONS ARE PLACED INTO ATOMIC ORBITALS BEGINNING
WITH THE LOWEST ENERGY ELECTRONS FOLLOWED BY SUCCESSIVELY HIGHER
-stable: electron occupy the lowest energy orbitals available: ground state
configuration of an atom by placing electron in the orbitals starting with the most
stable in energy and moving progressively upward.
-most stable: quantum number are not already assigned to another electron
Order of orbital filling
-An electron is spatially when it occupies one of the largest orbitals of the atom.
-Energetically when it occupies one of the least stable occupied orbitals of the atom
-electrons in less stable (higher energy) orbitals are thus more chemically active
than electrons in more stable orbitals.
- Accessible electrons: Valence electrons participate in chemical reactions
-Inaccessible electrons: Core electrons Orbital size increases and orbital stability decreases as the principal quantum
number n gets longer.
-Valence electrons are all those of highest principal quantum number plus those in
partially filled d and f orbitals.
5.3 Electron Configuration
Electron Configuartion: a complete specification of how