The Development of Atomic Models
• Democritus was a preSocratic Greek
philosopher (born
around 460 BC).
• Democritus was
originator of the belief
that all matter is made
up of various
imperishable, indivisible
elements which he
called "atomos", from
which we get the
English word atom.
•According to legend,
Democritus was supposed to
be mad because he laughed
at everything, and so he was
sent to Hippocrates to be
cured. Hippocrates pointed
out that he was not mad, but,
instead, had a happy
disposition. That is why
Democritus is sometimes
called the laughing
philosopher.
BB - Model
Dalton’s Model
More to come
Plum Pudding Model
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Proposed by J. J. Thomson
(1856 - 1940), the discoverer of
the electron in 1897.
The plum pudding model was
proposed in March, 1904 before
the discovery of the atomic
nucleus.
In this model, the atom is
composed of electrons
surrounded by a soup of positive
charge to balance the electron's
negative charge, like plums
surrounded by pudding. The
electrons were thought to be
positioned throughout the atom.
Electrons could move like letters
in alphabet soup
Instead of a soup, the atom was
also sometimes said to have had
a cloud of positive charge.
Thomson's model was compared
(though not by Thomson) to a
British treat called plum pudding,
hence the name. It has also been
called the chocolate chip cookie
model, but only by those who have
not read Thomson's original paper
Nuclear Model
• The Gold foil experiment
or the Rutherford
experiment was an
experiment done by
Ernest Rutherford (1871 1937) in 1909. This
experiment discovered the
nucleus.
• Led to the downfall of the
plum pudding model of
the atom.
• Alpha particles (positive
particles--Helium Nuclei)
were shot at gold foil.
• Particles passed through
the gold foil. A few shot
back.
• Conclusions:
1. Atom is mostly empty space
2. Dense center called the nucleus
3. Electrons were stuck surrounding
the nucleus.
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Planetary Model
• Introduced by Niels Bohr, a
Danish physicist (1885 1962), in 1913.
• Because of its simplicity,
the Bohr model is still
commonly taught to
introduce students to
quantum mechanics.
• The Bohr model depicts
the atom as a small,
positively charged nucleus
surrounded by waves of
electrons in orbit — similar
in structure to the solar
system, but with
electrostatic forces
providing attraction, rather
than gravity.
"The opposite of a correct statement is a
false statement. But the opposite of a
profound truth may well be another
profound truth." Niels Bohr
Quantum Mechanical Model
• Erwin Schrödinger (August
12, 1887 – January 4, 1961)
• An Austrian physicist,
achieved fame for his
contributions to quantum
mechanics, especially the
Schrödinger equation, for
which he received the Nobel
Prize in 1933.
• This model is based on
probability
• Where are you going to find
and electron 90% of the time.
• Atom is viewed as a fuzzy
cloud.
• Schrödinger equations create
electron clouds (orbitals) with
specific shapes.
Main Points For “Atoms” Video
• What is the key to understanding atomic structure?
• The discovery of what particle is associated with the
Crook’s Tube?
• What did Rutherford expect to happen in the gold foil
experiment?
• What was Rutherford’s genius?
• What conclusions did Rutherford draw from the Gold
foil experiment?
• How much smaller is the nucleus than the electron
cloud?
• What determines the shape of the electron cloud?
Dalton’s Atomic Theory
• Democritus - Greek philosopher who
first suggested atoms
• John Dalton (1766-1844)
• Studied ratios in which elements
combine
• Dalton put together the first atomic
theory
Dalton’s Atomic Theory
• All elements are composed of tiny indivisible particles called
atoms
• Atoms of the same element are identical. The atoms of any one
element are different from those of any other element
• Atoms of different elements can physically mix together or can
chemically combine with one another in simple-whole number
ratios to form compounds
• Chemical reactions occur when atoms are separated, joined, or
rearranged. Atoms of one element, however, are never changed
into atoms of another element as a result of a chemical reaction
Finding how many subatomic
particles for each atom
• Atomic Number - whole number on p.t.
• Gives the number of protons
• Atoms are electrically neutral; there,
positives equal negatives.
• Atomic number also equals number of
electrons
Finding how many subatomic
particles for each atom
• Mass number = protons plus neutrons
– Mass number = p+
+
no
• Mass number is not found on the
periodic table
• So, nO = mass number - p+
• If carbon has a mass number of 14, how
many e-’s, p+’s, and no’s does it have?
Symbols
• Mass number is in top left and atomic
number is in bottom left
9Be
• 16O
8
4
How many subatomic particles in each?
Oxygen Beryllium -
Isotopes
• Atoms with the same number of protons
but different numbers of neutrons
• Ex) Carbon 12 vs. Carbon 14
• These atoms have a different mass
• Chemically alike because still have the
same number of protons
Isotopes of Hydrogen
• Hydrogen -1 simply called hydrogen
• Hydrogen - 2 called deuterium
• Hydrogen - 3 called tritium
Development of AMUs
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Atomic Mass Units (AMUs)
Protons have a mass of 1 amu
1.67 x 10-24 g
Neutrons have a mass of 1 amu
Electrons have a mass of 0 amu
9.11 x 10-28 g
Atomic Mass
• The weighted average mass of the isotopes
in a naturally occurring sample of the element
• Don’t confuse with “mass number”
• To calculate atomic mass you need 3 pieces
of information
• 1. The number of stable isotopes
• 2.The mass of each isotope
• 3.The natural percent abundance of each
isotope
Atomic Mass
• Example Problem - Calculate the atomic mass for
element X. One isotope has a mass of 10 amus (10X)
and is 20% abundant. The other has a mass number
of 11 amus (11X) and an abundance of 80%.
• To solve: Multiply the mass number times the
abundance than add them together.
Atomic Mass
• 10 x 0.20 = 2.0
• 11 x 0.80 = 8.8
• Add 2.0 + 8.8 = 10.8
– The atomic mass of element X is 10.8
amus
Atomic Mass
• Your turn. Solve:
– What is the atomic mass of Element Z?
The isotopes are 16Z, 17Z, 18Z; with percent
abundances of 99.759, 0.037, 0.204.
Atomic Mass
• Answer
Atomic (Hotels) Theory
• Floors of the hotel are known as energy levels
• The period corresponds to the floor. The number of the
floor is called the Principle Quantum Number
• Energy levels are divided into sublevels. These are the
rooms of our atomic hotel. There are different types of
rooms.
• Atomic Orbital - Region of high probability of finding an
electron
• There are 4 types of rooms (orbitals). s, p, d, f rooms
• s - Spherical p - dumbbell d- clover leaf f - too
complex to describe.
• s - superior p - preferred d- desirable f- fair
– Sharp
principal
diffuse
fundamental
Atomic (Hotels) Theory
• Three Managers each with their own rule:
• 1. Aufbau principle
• Rooms closest to the basement are checked
out first
• Electrons enter orbitals of lowest energy first
• 2. Pauli Exclusion Principle - no more than 2
per room
• An atomic orbital contains a maximum of 2
electrons
Atomic (Hotels) Theory
• 3. Hund’s Rule
• Best rooms are check out first
• Similar rooms must have an occupant before pairing
up as roommates
• Roommates must have opposite spins to share a
room
• When electrons occupy orbitals of equal energy, one
electron enters each orbital until all the orbitals
contain one electron with parallel spins
S-orbitals of the 1st and 2nd energy levels.
P-orbital
2nd energy level
The anatomy of the periodic table
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Get out your periodic tables
Know where the following are on your periodic table (p.t)
Group or Family
Period
Group A (representative elements)
Group B
Metals
Nonmetals
Metalloids (Semimetals)
– Note - aluminum is not considered a metalloid
The anatomy of the periodic table
• Know where the following are on your periodic table
(p.t) continued
• Transition metals
• Inner transition metals
• Alkali metals
• Alkaline metals
• Halogens
• Noble gases
• Atomic Number
• Atomic Mass
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Focus Questions to upcoming video
8 Questions
What is the periodic law?
Who is Dmitri Mendeleev and what did he do?
How is today’s periodic table different from Mendeleev’s?
What characteristic is common among the noble gases
What are the names for vertical columns and horizontal
rows
• What characteristics are common amongst group 1A?
• What are the 2 most important things about the periodic
table?
• Why is fluorine the Tyrannosaurus rex of the periodic
table?
-1 0
10
Atoms
m
Nu cleus (protons an d neu trons)
Space occupied by electrons
Proton
Neu tron
10
-1 5
m
Periodic Table and Electron Configurations
• Build-up order given by position on periodic table; row
by row.
• Elements in same column will have the same outer
shell electron configuration.
The relation between orbital filling and the periodic table
Electron Configuration
• Orbitals have definite shapes and
orientations in space
(insert Fig 2.11 of text)
(if it will not all fit on one screen, put part
(a) on one screen and part (b) on the
next )
Orbital occupancy for the first 10 elements, H through Ne.
Trends in the
Periodic Table
Atomic radii of the maingroup and transition
elements.
Trend for atomic radii
• Left to right atoms get smaller
• Why?
– Increase in nuclear charge
– More protons and more electrons means greater
electrostatic attractions (stronger magnet)
• Top to bottom atoms get larger
• Why?
– Increase in energy levels (You are adding floors to your
hotel). Electrons are further from the nucleus
Atomic Radius
• Atomic radii actually
decrease across a row in
the periodic table. Due
to an increase in the
effective nuclear charge.
• Within each group
(vertical column), the
atomic radius tends to
increase with the period
number.
Atomic Radii for Main Group
Elements
Trend for Ion Size
• Ion is a charged atom.
• Metals lose electrons and nonmetals gain electrons to create
ions.
• Cations are pawsitive (positive) and Anions are negative.
• Cations are smaller than their corresponding atom. Why?
• Loss of electrons means the positive nucleus has a greater
attraction on the remaining electrons
• Anions are larger than their corresponding atom. Why?
• Gain of electrons means the nucleus has less attraction for the
electrons as well as the electrons are repulsing each other
causing an increase in the size of the electron clouds
Radii of ions
This is a “self-consistent” scale based
on O-2 = 1.40 (or 1.38) Å.
Ionic radii depend on the magnitude
of the charge of the ion and its
environment. (more later)
Positively charged ions are smaller
than their neutral analogues because
of increased Z*.
Negatively charged ions are larger
than their neutral analogues because
of decreased Z*.
Same periodic trends as atomic
radii for a given charge
Trend for ion size
• Decrease across a period then jumps in
size at nonmetals and continues to
decrease
• Increases on the way down a group as
you are adding energy levels (electrons
are farther from the nucleus)
Ionization energy
• The energy required to remove an
electron
First ionization
energies of the
main-group
elements
Trends in the
Periodic Table
Ionization Energy
• Ionization energy is a periodic property
Ionization energy
• In general, it increases across a row. Why?
• increasing attraction as the number of protons in the
nucleus increases (stronger magnet)
• it decreases going down a group. Why?
• Outer shell electrons are further from the nucleus so
less electrostatic attraction. Nucleus has less pull on
them.
• Shielding also plays a factor.
6) The trend across from left to right is
accounted for by a) the increasing nuclear
charge.
Electronegativity - This is the most
important trend to understand for this class.
• The tendency for an atom to attract electrons
when chemically bonded.
• Same trend as ionization energy.
– In general, it increases across a row. Why?
– increasing attraction as the number of protons in
the nucleus increases (stronger magnet)
– it decreases going down a group. Why?
– Outer shell electrons are further from the nucleus
so less electrostatic attraction. Nucleus has less
pull on them. Shielding also plays a factor.
Trends in three atomic properties
See chart in book for
summary
Check for understanding
• Which of the following atoms has the
largest atomic radii, ion size,
electronegativity, and ionization energy
• Na, Mg, K, Ca, S, Cl, Se, Br
Bonding
Ionic
Metallic
Covalent
Valence Electrons
• Atoms in a group behave similarly
because they have the same number of
valence electrons.
• Valence electrons - electrons in the
highest occupied energy level
• To find the number of valence electrons
just look at the group number
Lewis Electron Dot Structures
• Symbol of element with dots around it representing
valence electrons
• Example:
C
 2 electrons per side totaling 8
 No pairs until each room has an electron
Lewis Electron Dot Structures
• Example:
O
Unshared
Pair or Lone
Pair
• Pairs of electrons are adjacent not
across from one another. This will help
with identifying shape.
Octet Rule
• In forming compounds, atoms tend to achieve the
noble gas electron configuration.
• They lose, gain, or share electrons with another atom
to achieve 8.
• Metals lose electrons leaving a complete octet in the
lower energy level
• Nonmetals gain electrons to fill the energy level to
achieve 8.
Electron Configurations of Ions
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Na - 1s22s22p63s1 Na1+ - 1s22s22p6
O - 1s22s22p4
O2- - 1s22s22p6
Write the following on your periodic table.
Group 1A - 1+
- loses 1 electron
Group 2A - 2+
- loses 2 electrons
Group 3A - 3+
- loses 3 electrons
Group 4A - Depends on the atom
Group 5A - 3- gains 3 electrons
Group 6A - 2- gains 2 electrons
Group 7A - 1- gains 1 electron
Group 8A (0) - does not form ions
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VSEPR Theory
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Valence Shell Electron Pair Repulsion Theory
“Electron pairs around atoms tend to be as far apart as possible.”
Similar charges (I.e., negative charges from electrons) tend to repel each other
and want to be spaced apart at maximum angles.
Used to predict molecular geometries
Bond angles
– Angles between bonds
– Spacing apart as far as possible
Lone pairs of electrons will repel bonded atoms a bit more than expected
toward each other around the central atom
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Covalent Bonding
Polar Bonds and Molecules
Covalent Bonding
-- Polar Bonds and Molecules -Bond Polarity
• “The Tug of War”
– The pairs of electrons that are bonds between atoms are pulled
between the nuclei of the atoms in a bond.
– The electronegativities of the atoms determines who is winning
Yet; there is no winner. The tug of war never ends.
• Classifications for Bonds
– Nonpolar covalent
• When atoms pull the bond equally
• Happens with two atoms of equal electronegativity, most often
using the same atoms
• Examples: H2, O2, N2
– Polar covalent
• When atoms pull the bond unequally
• Happens with two atoms of different electronegativities
• Example: HCl, HF, NH
Covalent Bonding
-- Polar Bonds and Molecules -Bond Polarity
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•
In a polar molecule, one end of the molecule is slightly more electronegative
than the other atom, resulting in one atom being slightly negative (-) because
of higher electronegativitiy, and the other atom being slightly positive (+)
because of lower electronegativity.
 is known as a partial charge since it is much less than 1+ or 1- charge.
Covalent Bonding
-- Polar Bonds and Molecules -Bond Polarity
• Electronegativities and Bond Types
– H: 2.1 Cl: 3.0 Since hydrogen is less, it will have the positive
partial charge while chlorine has the negative partial charge.
– 3.0 – 2.1 = 0.9 HCl is polar covalent.
0.0 – 0.1 difference
Nonpolar covalent bond
H – H (0.0 difference)
0.1 – 1.7 difference
Polar covalent bond
H – Cl (0.9 difference)
1.7 + difference
Ionic bond
Na+Cl- (2.1 difference)
Covalent Bonding
-- Polar Bonds and Molecules -Polar Molecules
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•
Dipole
– Molecule that has two poles
– Example: HCl from the previous page
Polar vs. Nonpolar
H2O and CO2
Both have 3 atoms; yet,
One is polar and one is
nonpolar.
Why?
Structure (with bond
polarity) determines the
molecules polarity.
3 video clips coming up
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• Get out your bonding sheet of chemistry
• Yes, we are going to discuss more on bonding so try
and hold your enthusiasm as rioting is not tolerated
and things need to be accomplished.
• Bonding appreciates your cooperation and will sign
autographs at the end of the period.
• Thank you.
Intermolecular Attractions
Attractions Between Molecules
• van der Waals forces
– Two types: dispersion forces and dipole interactions
• Dispersion forces
– Weakest of all molecular interactions
– Caused by movement of electrons
– Occurs in the BrINClHOF’s
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Intermolecular Attractions
Attractions Between Molecules
2nd van der Waals force
• Dipole interactions
• Occurs when polar molecules are attracted to one
another
• Partial charge (+) of one polar molecule is
attracted to the opposite partial charge (-) of
another molecule
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Intermolecular Attractions
attractions between molecules
• Hydrogen bonding
– Hydrogen covalently bonded to a very electronegative atom is also
weakly bonded to an unshared electron pair of another
electronegative atom
– Example: water
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Short Lab - No lab report
• In the back in a tray is a micropipet and a penny.
Place 1 drop of water on the penny and then take a
guess as to how many drops of water you can fit on a
penny. Write your guess on the board at the front of
the room. Place your name next to your guess. Then
count the drops of water you fit on the penny before it
overflows. Record this count on the board next to
your guess. As your water drop grows, watch it from
the side. Clean up and have a seat at your desk.
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Gases
• There are four variables that affect a
gas.
• 1. Pressure
• 2. Volume
• 3.Temperature
• 4. Number of molecules
The variables
• Pressure units - there are many units for pressure.
– kPa - kilopascal (101.3)
– Atm - atmospheres (1 )
– mm Hg - millimeters of Hg (760) - torr
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Volume is measured in Liters
Temperature is in Kelvin
K = oC + 273 or oC = K - 273
If the Kelvin temperature doubles the K.E. doubles.
The pressure-volume relationship
Boyle’s Law
• Pressure and volume are inversely
related.
• One goes up the other goes down
• P1 x V1 = P2 x V2
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Boyle’s Law
sample problem
• A high-altitude balloon contains 30.0 L
of helium gas at 103 kPa. What is the
volume when the balloon rises to an
altitude where the pressure is only 25.0
kPa?
• 103 kPa x 30.0 L = 25.0 kPa x V2
• V2 = 124 L
Boyle’s Law
sample problem
• Your turn
• Your birthday balloon travels with you
from Lincoln to Denver. The balloons
volume is 4.0 L with an atmospheric
pressure of 101.3 kPa. You arrive in
Denver where the atmospheric pressure
is 90.0 kPa. What is the new volume of
your balloon?
Boyle’s Law
sample problem
• Answer
The temperature-volume relationship
Charles’s Law
• Volume and temperature have a direct
relationship.
• One goes up the other goes up
– One goes down the other goes down
V1
T1
=
V2
T2
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Charles’s Law
sample problem
• A balloon inflated in a room at 24oC has
a volume of 4.00 L. The balloon is then
heated to a temperature of 58oC. What
is the new volume if the pressure
remains constant?
4.00L = V2
297 K 331 K
V2 = 4.46 L
Charles’s Law
sample problem
• Your turn
• A balloon is inflated in a room at 24oC
and has a volume of 4.00 L. The balloon
is placed in a freezer and then removed
the volume is now 3.25 L. What was the
temperature of the freezer in Celsius?
Charles’s Law
sample problem
• Answer
The Temperature-Pressure Relationship
Gay-Lussac’s Law
• The pressure of a gas is directly
proportional to the temperature of a gas
• Temperature goes up; pressure goes up
P1 = P2
T1 T2
Gay-Lussac’s Law
example problems
• The gas left in a used aerosol can is at
a pressure of 103 kPa at 25oC. If the
can is thrown into a fire, what is the
pressure of the gas when it reaches
928oC?
103 kPa = P2
298 K
1201 K
P2 = 415 kPa
Gay-Lussac’s Law
example problems
• Your turn
• A container of propane has a pressure
of 108.6 kPa at a morning temperature
15oC. By mid afternoon the temperature
has reached 32oC. What is the pressure
inside the propane tank?
Gay-Lussac’s Law
example problems
• Answer
The combined gas law
P1 x V 1
T1
=
P2 x V 2
T2
The combined gas law
example problem
• The volume of a gas-filled balloon is
30.0 L at 40oC and 153 kPa. What
volume will the balloon have at STP?
153 kPa x 30.0 L = 101.3 kPa x V2
313 K
273 K
V2 = 39.5 L
The combined gas law
example problem
• Your turn
• A gas-filled balloon is 25.0 L at 35oC
and 145 kPa. What is the temperature if
the volume increases to 28.0 L and a
pressure of 152 kPa?
The combined gas law
example problem
• Answer
The Ideal Gas Law
PV=nRT
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•
P = Pressure
V = Volume
n = Number of Moles
R = ideal gas constant
T = temperature in Kelvin
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R -The ideal gas constant
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Depends on unit of pressure
0.0821 L . Atm / K . mol
62.4 L . mmHg / K . mol (torr is mm Hg)
8.31 L . kPa / K . mol
Ideal Gas Law
example problem
• Calculate the pressure of 1.65 g of helium gas at 16.0oC and
occupying a volume of 3.25 L?
• You will need g to moles and Celsius to Kelvin:
• 1.65 g He 1 mole He
•
4.0 g He = 0.413 mol He
• K = oC + 273 ; 16. 0 + 273 = 289 K
• For this problem you will need to pick an R value. For this
problem I will choose to use the R value containing kPa. I picked
it. You can’t do anything about it. So; just try and stop me.
• Plug and Chug baby, get ‘R done. Do it. Come on I dare ya.
• Get it - ‘R as in ideal gas constant
Ideal Gas Law
example problem
• P x 3.25 L = 0.413 mol x 8.31 kPa . L x 289 K
•
mol . K
• Do the algebra and solve; if you do it right, guess
what? You get the answer right. Neat concept, huh?
Maybe your mommy will give you a cookie.
• = 305 kPa
• Your turn
• How many moles of gas are present in a sample of
Argon at 58oC with a volume of 275 mL and a
pressure of 0.987 atm.
Ideal Gas Law
example problem
• Answer
• 0.987 atm x 0.275 L = n x 0.0821 L . Atm x 331K
•
mol . K
• Do the dew; oops, I mean the algebra and presto; the answer
with the correct number of sig figs is..
• Do you know how to keep a so called chem student in
suspense?
• Do ya?
• Do ya?
• = 0.00999 mol Ar
• Congrats - you can plug and chug.
• Bye Bye now.
Surface Tension
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Intermolecular Forces Bulk
and Surface
Phase Changes
Energy Changes Accompanying Phase Changes
• All phase changes are possible under the right
conditions (e.g. water sublimes when snow
disappears without forming puddles).
• The sequence
heat solid  melt  heat liquid  boil  heat gas
is endothermic.
• The sequence
• cool gas  condense  cool liquid  freeze  cool solid
is exothermic.
Phase Changes
Heating Curves
Heating Curve Illustrated
Vapor Pressure
Explaining Vapor Pressure on
the Molecular Level
• Dynamic Equilibrium: the point
when as many molecules
escape the surface as strike
the surface.
• Vapor pressure is the pressure
exerted when the liquid and
vapor are in dynamic
equilibrium.
Vapor Pressure
Volatility, Vapor Pressure, and
Temperature
• If equilibrium is never established then the
liquid evaporates.
• Volatile substances evaporate rapidly.
• The higher the temperature, the higher the
average kinetic energy, the faster the liquid
evaporates.
Liquid Evaporates when no Equilibrium is
Established
Vapor Pressure
Vapor Pressure and Boiling Point
• Liquids boil when the external pressure equals the vapor
pressure.
• Temperature of boiling point increases as pressure
increases.
• Two ways to get a liquid to boil: increase temperature or
decrease pressure.
– Pressure cookers operate at high pressure. At high
pressure the boiling point of water is higher than at 1 atm.
Therefore, there is a higher temperature at which the food
is cooked, reducing the cooking time required.
• Normal boiling point is the boiling point at 760 mmHg (1
atm).
Thermochemistry
-- The Flow of Energy: Heat --
Thermochemistry: the study of heat
changes in chemical reactions
Chemical potential energy: energy stored
within the structural units of chemical
substances
Thermochemistry
-- The Flow of Energy: Heat --
Chemical System Types
System type
Endothermic
Exothermic
Description
System absorbing heat
from the surroundings
System releasing heat to
the surroundings
q (change
in heat)
q > 0
q < 0
Thermochemistry
-- The Flow of Energy: Heat --
Law of Conservation of Energy:
In any chemical or physical
process, energy is neither
created nor destroyed
Thermochemistry
-- The Flow of Energy: Heat -The calorie
• Expressed as a c (lower case)
• Quantity of heat needed to raise the temperature
of 1 g of pure water 1C
Calorie
• Expressed as a C (upper case)
• Dietary Calorie
• 1 Calorie = 1 kilocalorie = 1000 calories
Thermochemistry
-- An Intro Video --
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Thermochemistry
-- The Flow of Energy: Heat -Joule
• SI unit of heat and energy
• Raises the temperature of 1 g of pure water 0.2390C
• 4.184 J = 1 cal
Heat Capacity
• Amount of heat needed to increase the temperature of an object
exactly 1C
• Will change depending on the mass and chemical composition
Specific Heat
Quantity of heat needed to raise the temperature of
1g of substance 1oC
Thermochemistry
-- The Flow of Energy: Heat --
Specific Heat Capacity
Heat (q)
Mass (m)
specific heat capacity (C)
change in temperature (T)
q = mC T
Thermochemistry
-- The Flow of Energy: Heat -Example:
How many kilojoules of heat are absorbed when 1.00 L of water is
heated from 18C to 85C?
Solution:
q = mCT
q = 1000g x 4.18 J x 67oC
goC
q = 2.8E5 J 1 KJ
1000 J
= 280 KJ
Thermochemistry
-- The Flow of Energy: Heat -Example:
A chunk of silver has a heat capacity of 42.8 J/C. If the silver
has a mass of 181 g, calculate the specific heat of silver.
Solution:
q = mCT
42.8 J = 181g x C x 1OC
C = 0.236 J/goC
Thermochemistry
-- Measuring and Expressing Heat Changes -Your Turn:
The temperature of a piece of copper with a mass of
95.4 g increases from 20.0oC to 43.0oC when the
metal absorbs 849 J of heat. What is the specific
heat of copper?
Thermochemistry
-- Measuring and Expressing Heat Changes --
Calorimeter
Properties of acids and Bases
• Taste
– Acids taste sour ex - lemons
– Bases taste bitter ex - soap
• Feel
– Acids feel like water; but have you ever
gotten fruit juice on a canker sore or cut
– Bases feel slippery ex - soap and water
Properties of acids and bases
• Reaction with metals
– Acids - Hydrogen gas is produced when
reacted with certain metals
– Bases - typically don’t react with metals
• Both are electrolytes
• React to form salt and water
• Milk of Magnesia (Magnesium
Hydroxide) is a base used to treat
excess stomach acid problems.
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Definitions of Acids and Bases
• Arrhenius Acid/Base - focused on
products
– HCl + H2O --> H3O+ + Cl– NaOH + H2O --> Na+(aq) + OH-(aq)
• Acids form H+’s and Bases form OH-’s
Definitions of Acids and Bases
• Bronsted-Lowery Acid/Base - focused
on what happens during formation
– HCl + H2O --> H3O+(aq) + Cl-(aq)
• Acid - substance that donates a proton
• Base - substance accepts a proton
• In the above example, what is the base
and what is the acid?
• How about this example?
– NH3 + H2O --> NH4+(aq) + OH-(aq)
Definitions of Acids and Bases
• Lewis acid/base
– H+ + OH- ---> H2O
• Acid - a substance that can accept a
pair of electrons to form a covalent bond
• Base - a substance that donates a pair
of electrons to form a covalent bond
• What is the Lewis acid and base?
– AlCl3 + Cl- --> AlCl4-
Problem
• Write an equation for the ionization of
nitric acid and explain how it fits each
definition?
• HNO3 --> H+ + NO3- or
• HNO3 --> H3O+ + NO3- or
• HNO3 + H2O --> H3O+ + NO3-
Hydrogen Ions from Water
This is the “bases” to start understanding pH
• Water is considered neutral
• Collision between water molecules can cause
a hydrogen ion to transfer from one molecule
to another.
H2O
H2 O
H3O+
hydronium ion
OHhydroxide ion
Self-Ionization of Water
Water self ionizes to the concentration of 1.0 x 10-7 mol/L.
When the concentration of each ion equals 1.0 x 10-7 mol/L
the solution is said to be neutral
Therefore, since water is considered neutral the
concentrations of the ions can be calculated through the
Ion product constant.
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The Ion-Product Constant
Kw
• Notation
– [H+] - concentration of hydrogen ions
• Or hydronium ions
– [OH-] - concentration of hydroxide ions
• When [H+] and [OH-] are multiplied we get the ionproduct constant.
• Kw = [H+] x [OH-] = 1.0 x 10-14 (mol/L)2 or M2
• This is an inverse relationship.
– One goes up, the other goes down.
The Ion-Product Constant
Kw example problem
• If [H+] = 1.0 x 10-5 mol/L, is the solution
acidic, basic, or neutral? What is the
[OH-] of this solution?
• Answer
– Acidic - the [H+] is greater than 1.0 x 10-7
mol/L
– 1.0 x 10-5 mol/L x [OH-] = 1.0 x 10-14 M2
– [OH-] = 1.0 x 10-9 mol/L
The Ion-Product Constant
Kw example problem
• If [OH-] = 2.8 x 10-8 mol/L, is the solution
acidic, basic, or neutral? What is the [H+] of
this solution?
• Answer
The pH concept
• [H+] is cumbersome so the pH scale
was created.
• pH is the negative logarithm of the
hydrogen-ion concentration.
• pH = -log[H+]
Sample pH problems
1 of 3
• The hydrogen-ion concentration of a
solution is 2.7 x 10-10 mol/L. What is the
pH of the solution?
• Answer
Sample pH problems
2 of 3
• The pH of a solution is 6.8. What is the
[H+]?
• Answer
Sample pH problems
3 of 3
• What is the pH of a solution if the [OH-]
= 4.0 x 10-11 mol/L
• Answer
Naming Acids
And writing formulas
General Form - HX (X is an
anion or polyatomic ion)
• Rules - 3 of them
• 1. When anion ends in -ide, acid name
begins with hydro and -ide is changed
to -ic with the word acid
• Ex - HCl is hydrochloric acid
• Try - H2S
• Rule 2 • When anion ends in -ite, ending changes to ous with the word acid. (No hydro)
• Name these - H2SO3 , HNO2
• Rule 3 • When anion ends in -ate, ending changes to ic with the word acid
• Name these - HNO3 , HC2H3O2
Work backwards to get the
formula
• Write the formula for the following:
– Chloric acid
– Hydrobromic acid
– Phosphorous Acid
• Don’t confuse phosphorous and phosphorus
• Answers
Neutralization Reaction
• Base and acid react to produce a salt
and water
Titration
• A lab technique where a neutralization
reaction is performed to determine the
concentration of an unknown.
The anatomy of a titration:
• Standard solution - solution of known
concentration.
• End Point - the point at which neutralization
is achieved
• Indicator - chemical that changes
color with a change in pH.
We will be using phenolphthalein.
Clear in an acid pink in a base
You want a light pink
• Buret - Measurement device
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Titration Calculations
4 steps
•
•
•
•
Start with the balanced equation
Find the moles in the standard solution
Set up ratio to find moles of unknown
Find molarity(mol/L) or volume
Titration Calculations
example problem
• A 25.75 ml solution of H2SO4 is neutralized by
18.23 ml of 1.0 M NaOH. What is the
concentration of H2SO4?
• H2SO4 + 2 NaOH --> Na2SO4 + 2 H2O
• 0.01823 L NaOH 1.0 mol NaOH 1 mol H2SO4
1 L NaOH
• = 0.35 mol/L H2SO4
2 mol NaOH 0.02575L H2SO4
Titration Problems
your turn
• What is the molarity of phosphoric acid if 15.0
mL of the solution is neutralized by 8.5 mL of
0.15 M NaOH?
Titration Problems
1 more your turn
• How many milliliters of 0.45 M hydrochloric acid
must be added to 25.0 mL of 0.15 M NaOH?