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Transcript
Atoms, Molecules, and Ions
Chapter 2
Naming Ionic Compounds
• Positive ion is always named first
• Negative ion is always named second
Tips for memorizing:
1. Learn the “ate” ion
2. The “ite” has one less oxygen
3. The prefix “per” is used with “ate” to
indicate one more oxygen
4. The prefix “hypo” is used with “ite”
to indicate one less oxygen
• Metals with more than one oxidation
state (transition metals) must have a
roman numeral to indicate the
oxidation state
• Monatomic anions end in “ide”
Fundamental Chemical Laws
Law of Conservation of Mass
• Mass of reactants equals the mass of the
products
Law of Definite Proportion
• Compounds have an unchanging chemical
formula
Law of Multiple Proportions
• Sometimes two elements can come
together in more than one way, forming
compounds with similar, though not
identical formulas
Dalton’s Atomic Theory
1. Each element is composed of Atoms
2. All atoms of a given element are
identical; atoms of different elements
have different properties
3. Atoms of an element are not changed
into different types of atoms by chemical
reactions; atoms are neither created nor
destroyed in chemical reactions
4. Compounds are formed when atoms
of more than one element combine;
a given compound always has the
same relative number and kinds of
atoms.
Discovery of Atomic Structure
1.
2.
3.
4.
JJ Thomson and the Electron
Robert Millikan’s Oil Drop
Radioactivity
The Nuclear Atom
JJ Thomson’s Cathode Rays
JJ Thomson and Cathode Rays
1. Determined the charge to mass ratio
of the electron
2. Reasoned that all atoms must
contain electrons
3. Reasoned that all atoms must
contain positive charges
Millikan’s Oil Drop
Milikan’s Oil Drop
1. Oil drop experiments determined
charge of an electron
2. Charge info and Thomson’s
charge/mass ratio allowed him to
determine the mass of an electron
at (9.11 x10-28 g)
Behavior of radiation in an Electric Field
Radioactivity
1. Gamma (g) rays – high energy light
2. Beta (b) rays – high speed electrons
3. Alpha (a) particles – nuclear particle
with a +2 charge
Rutherford’s Gold Foil Experiment
The Nuclear Atom
Rutherford’s Metal Foil Experiment
• Most alpha particles pass straight
through metal foil
• Some particles were greatly deflected
Why do you think? What could this prove?
• Must have been deflected by a
small, central, positive nucleus
where the atom’s mass was
concentrated
Modern View of the Atom
Nucleus:
Small size, high density
Contains:
1. Protons
• Positively charged
2. Neutrons
• No charge
Electrons
1. Negatively charge
2. Source of element’s reactivity
3. Provide most of the atomic volume
Subatomic particle relationships
• Atomic Number
Number of protons
• Mass Number
Number of protons + number of neutrons
• Isotopes
Atoms with the same number of protons (same
element), but different numbers of neutrons
(mass numbers)
If I change the number of protons
in an atom…
• I get a new element
If I change the number of electrons
in an atom…
• I get an ion
If I change the number of neutrons
in an atoms…
• I get an isotope
Symbols for the Elements
Atomic Weights
• Atomic Mass Scale: determined by
assigning a mass of exactly 12 amu to
an atom of the 12C isotope of carbon.
• Average Atomic Masses: most
elements exist as mixtures of isotopes
so avg. atomic masses is determined
by using the masses of its various
isotopes and their relative abundances.
Mass Spectrometer
• Instrument used to measure the precise
masses and relative amounts of atomic
and molecular ions
The Periodic Table
• Horizontal Row = “period” or “series”
• Vertical Column = “group” or “family”
Group 1A – Alkali metals
Group 2A – Alkaline earth metals
Group 7A – Halogens
Group 8A – Noble gases
Group B – Transition Metals
Metals, Metalloids, Nonmetals??
METALS
a)
b)
c)
d)
e)
Left side of periodic table
High electrical conductivity
High Luster (very shiny when clean)
Ductile (drawn into wires)
Solid @ room temperature (except Hg)
NON-METALS
a)
b)
c)
d)
Upper right corner of periodic table
Non-lustrous (not shiny)
Poor conductors of electricity
Some are gases at room temp.(O,N)
--others brittle solids (S)
Semi-Metals (metalloids)
• Divides metals from nonmetals
• properties intermediate between metals and
nonmetals
• Ex: Si, Ge, ...
• “stair case” on right
Side of periodic table
Valence Electrons
• Electrons in highest occupied
principle energy level
• Electrons available for bonding
• For Representative Elements:
–Number of valence electrons in
an atom = the group number of
that atom
has one valence electron.
has two valence electrons.
has three valence electrons.
And So On….
has seven valence electrons.
has eight valence electrons.
diagrams that show the valence
electrons as dots.
Since only the valence electrons are
involved in bonding, this is very
useful in predicting which atoms will
combine chemically.
6 3
4
7
X
2
1
5 8
The X represents the symbol for the
element.
The dots are placed around the
symbol in the order shown above.
The Octet Rule
• When forming compounds, atoms tend to
achieve the electron configuration of a
noble gas. (all 8 electrons in outer shell)
• Done by losing, gaining, or sharing electrons
Molecules and Molecular Compounds
Covalent Bonding:
• Sharing of electrons
• Typically between nonmetals and/or
large molecules
Diatomic Molecule –
a form of covalent bonding in which
two identical atoms are bonded
together (H2,N2,O2,Cl2,F2,Br2,I2)
Mono-
1
• EXAMPLES:
Di-
2
Tri-
3
• CO2
Tetra-
4
Penta-
5
Hexa-
6
Hepta-
7
Octa-
8
Nona-
9
Deca-
10
• Carbon Dioxide
• P4O6
• Tetraphosphorus
Hexoxide
Formulas for Nonmetal Compounds
(Covalent Compounds)
1. Write correct symbols for the
elements.
2. Use prefixes to add the correct
subscript.
 Example: dinitrogen tetroxide
 N2O4
Representing Molecules
• Chemical Formulas: H2O , CH4
• Structural Formulas:
– Bonds represented by lines
– Ball and Stick
– Space Filling
• Empirical Formulas: only gives the
relative number of atoms of each
type in a molecule (smallest possible
whole number ratio of atoms)
Ions and Ionic Compounds
• Cation:
Positive ions formed by the loss of
electrons
• Anion:
Negative ions formed by gaining
electrons
Ionic Bonding
1. Formed by attraction between
oppositely charge ions (electrons are
transferred)
2. Forms ionic solids (salts)
3. Can be monatomic (one atom) or
polyatomic (more than one atom)
4. Typically between a metal and a
nonmetal
Something to Ponder…
• How does a recipe that calls for 4 cups
of tomato sauce and 1 teaspoon of salt
and the following formula CaCl2 relate
to one another?
• 1 atom of Calcium for every 2 atoms of
Chloride
Naming Compounds
(Ionic Compounds)
1. Name the cation
 If it’s a transition metal, give it roman
numerals for its correct charge
2. Name the anion
•
•
•
•
•
•
•
•
Examples:
NaBr =
Sodium Bromide
CuO =
Copper Oxide ??
Copper (II) Oxide
Na2SO4 =
Sodium Sulfate
Name the following with yourIron
Group:
(II) Bromide
1. Na2O
9. FeBr2
10. Cu2O
11.CaCl2
2. LiBr
3. SnO2 Tin (IV) Oxide
4. Cu3P2
Copper (II) Phosphide
5. Sr3P2
12.AlP
6. BaO
13.FeCl3
7. PbBr4 Lead (IV) Bromide
14.MgBr2
8. Cr2O3
15.AlCl3
Chromium (III) Oxide
Writing Formulas
(Ionic Compounds)
1. Write the Symbol and Charge of the
cation
2. Write the Symbol and Charge of the
anion
3. Add needed subscripts to balance the
charges
 Find least common multiple (hard way)
 Use crisscross method (easy way)
3+
Fe
2O
Find the least common multiple.
In this case it is 6.
You must have 2 Fe and 3 O to
balance the charges.
Fe2O3
3+
Fe
2O
The easy way to do this is to “crisscross” the
charges and make them subscripts.
Fe2O3
Note: Always remember to reduce the
subscripts to lowest terms.
Try a Couple:
•
•
•
•
•
•
•
Calcium Nitride
Ca3N2
Lithium Oxide
Li2O
Barium Sulfide
BaS
Why isn’t it Ba2S2 ??
Formulas w/ Polyatomic Ions:
1. Same steps as with monatomic ions
2. must use parenthesis to enclose
polyatomic ion if more than one is
needed
 Example: Calcium Nitrate
 Ca 2+ NO3 1 Ca(NO3)2
Naming Acids
• Binary Acids (2 elements – hydrogen
+ one other)
– Prefix “Hydro” + root of 2nd element
+ “ic”
• Oxyacids
– If the acid contains an anion whose
name ends in “ate”: use the root of
anion name and an “ic” ending
– If acid contains an anion whose name
ends in “ite”: use the root of the
anion name and an “ous” ending
Electronic Structure of Atoms
Chapter 6
Light
• Made up of
electromagnetic
radiation.
• Waves of electric
and magnetic fields
at right angles to
each other.
Parts of a wave
Wavelength
l
Frequency (n) = number of cycles in one
second
Measured in hertz
1 hertz = 1
Frequency = n
Kinds of EM waves
• There are many different EM waves
• different l and n
• Visible Light is only the part our eyes can
detect. (colors of the rainbow)
• Greater wavelength means, smaller
frequency
Gamma
Rays
MicroX-Rays
UV
Radio
Infrared
wave
The speed of light, c
• in a vacuum is 2.998 x 108 m/s
• c = 3.0 x 108 m/s
• c = ln
Examples
What is the wavelength of light
with a frequency 5.89 x 1014 Hz?
l =
c
v
=
3.0 x 108 m/s
5.89 x 1014 Hz
= 5.09 x 10-7 m = 509 nm
(green light)
What is the frequency of blue light
with a wavelength of 484 nm?
v=
c
l
=
3.0 x 108 m/s
484 x 109 m
= 6.20 x 1014 Hz
Planck and the Quantum Theory
• Energy is gained or lost in whole number multiples
(n) of the quantity hv. (Similar to energy required
to go up stairs vs. up a ramp)
• Frequency = v
• Planck’s constant = h = 6.63 x 10-34 J-s
• Planck found discovered that Energy is transferred
to matter in “energy packets” called a quantum
(hv)
DE = nhn
Einstein, the Photoelectric Effect,
and Photons
• EM radiation is quantized a stream of
particles -- “photons”
• Ephoton = hn = hc/l
• Combine this with E = mc2
• You get the apparent mass of a photon.
m = h / (lc)
Is light a Wave or
does it consist of particles?
• Both…
• Macroscopically like a wave,
• But consists of a collection of photons
that we only see at the atomic level.
• called The Wave-Particle Duality
(Like describing an entire beach and then
beginning to examine the grains of
sand.)
Examples
• Calculate the energy of one photon of
yellow light whose wavelength is
589nm
1. Find the frequency
• 5.09 x 1014 s-1
2. Then use Plank’s equation to find E
• 3.37 x 10-19 J
Matter as a wave
• Using the velocity (v) instead of the
frequency (n) we get:
• De Broglie’s equation l = h/mv
• Can calculate the wavelength of an
object.
Line Spectra
• Spectrum = the range of frequencies
present in light
• Continuous Spectrum = contains all
wavelengths of light. (white light… can
be broken down into “rainbow”)
• Line Spectrum = contains only specific
wavelengths of light.
Hydrogen spectrum
• Emission spectrum because these are the colors
it gives off or emits.
• Called a bright line emission spectrum.
• There are just a few discrete lines showing
656 nm
434 nm
410 nm
486 nm
Bright Line Spectra
• Excited electrons return to lower NRG states
• NRG is emitted in the form of a photon of definite
wavelength.
• Definite change in energy corresponds to:
– Definite frequency
– Definite wavelength
• Use DE = hn = hc / l
• Only certain energies are possible within any atom.
Niels Bohr
• Developed the Quantum Model
• Described the atom like a solar system
• Electrons attracted to (+) nucleus
because of their (-) charge
• Electrons didn’t fall into nucleus because
they were moving around
Bohr’s atom
• Found only certain NRGs were allowed;
called them NRG levels.
• Putting NRG into atom moves electron
away from the nucleus (ground state 
excited state)
• When e- returns to ground state, it
gives off light of a certain NRG
The Bohr Atom
n=4
n=3
n=2
n=1
Available NRG levels
E = -2.178 x 10-18 J (Z2 / n2 )
• n = quantum number (NRG level)
• Z = nuclear charge (+1 for Hydrogen)
• J = energy in joules
• The more negative the NRG is, the more
stable the atom will be.
change in Energy
• When the electron moves from one
energy level to another:
• DE = Efinal - Einitial
DE = -2.178 x 10-18J [(1/ nf2)–
(1/ ni2)]
l = hc / DE
Shortcomings of Bohr Model
• Only works for Hydrogen atoms
• Electrons don’t move in circular orbits
• The quantization of energy is right, but not
because they are circling like planets
• Questions Bohr couldn’t answer:
Why are e- confined to only certain energy levels?
Why don’t e- eventually spiral and crash into the
nucleus?
The Quantum Mechanical Model
• New approach that viewed electron
as a standing wave of NRG
• Standing waves don’t propagate
through space
• Standing waves are fixed at both ends
(similar to vibrations of a stringed
instrument)
What’s possible?
• You can only have a standing wave if you have
complete waves.
• There are only certain allowed waves.
• In the atom there are certain allowed waves
called electrons.
• 1925 Erwin Schroedinger described the wave
function of the electron. “The Schroedinger
Equation”
• Much math but what is important are the
solutions.
Schroedinger’s Equation

2x2
22
•
•
•
•
•
•

2y2
22

2z2
22


82m
h2

(E  V)  = 0
The wave function,  is a F(x, y, z)
Solutions to the equation are called orbitals.
These are not Bohr orbits.
Each solution is tied to a certain energy.
These are the energy levels.
Many strange and seemingly impossible
behaviors occur when the electron is treated as a
wave!
Orbitals
• Orbitals are not circular orbits for
electrons
• Orbitals are areas of probability for
locating electrons
There is a limit to what we can
know…
• about how the electron is moving or how it
gets from one energy level to another.
• about both the position and the momentum
of an object.
• The Heisenberg Uncertainty Principle “we cannot know the exact location and
exact momentum of an electron at the
same time.”
Quantum Mechanical Model and
Quantum Numbers
• Note: A quantum mechanical orbital is
not the same as a Bohr orbit because
the motion of the electron in an atom
cannot be precisely measured or
tracked. (Heisenberg uncertainty Principle)
• There are 4 quantum numbers to describe
the “location” of an electron. (sort of like
how a zip code works)
Principal Quantum Number (n)
• Indicates probable distance from the
nucleus (old Bohr orbitals)
• Gives the size and energy of the orbital
• Has integer values >0
• According to the periodic table, what
would the highest principal quantum
number be?
Angular Momentum
Quantum (l )
• Gives the shape of the orbital (more detail to
come)
• Integral values from 0 to (n-1) for each principal
quantum number (n)
Value of l
0 1 2 3 4
Letter used for shape* s p d f g
*letters s, p, d, f come from the words sharp, principal, di
and fundamental, which were used to describe certain fea
of spectra before quantum mechanics was developed.
Magnetic Quantum Number (ml
)
• Relates to the orientation of the orbital
in space relative to the other orbitals.
(It tells you if the orbital will be on the
x, y or z axis.)
• Integral values from l to –l including 0.
n
l
1
2
3
4
0
0
1
0
1
Orbital
designation
1s
2s
2p
3s
3p
ml
0
0
-1, 0, 1
0
-1, 0, 1
# of
orbitals
1
1
3
1
3
2
0
1
2
3
3d
4s
4p
4d
4f
-2, -1, 0, 1, 2
0
-1, 0, 1
-2, -1, 0, 1, 2
-3, -2, -1, 0, 1, 2, 3
5
1
3
5
7
Important Observations
1. The shell w/ quantum #n will have exactly n
subshells.
2. Each subshell has a specific number of orbitals.
Each orbital corresponds to a different allowed
value of ml. For a given value of l, there are
2l + 1 allowed values of ml.
3. The total number of orbitals in a shell is n2. The
resulting number of orbitals for the shells – 1, 4,
9, 16 – is related to a pattern seen in the
periodic table… We see the number of elements
in the table – 2, 8, 18, 32 – equal twice these
numbers…
S orbitals
n=1 n=2
n=3
P orbitals
At another energy level the solutions are
“dumbell” shaped.
There are 3 possible solutions for this energy leve
P Orbitals
All 3 p orbitals may exist at the same time.
d orbitals
At another energy we get “flower” shaped orbitals for a solution.
All 5 may exist
at the same
time
F orbitals
And finally, at another energy, 7 f orbitals are the solution.
Orbital Energies
• All orbitals with the same value of n
have the same energy
• The lowest energy state is called the
“ground state”
• When the atom absorbs energy,
electrons may move to higher energy
orbitals – “excited state”
Electron Spin
Quantum Number (ms )
• An individual orbital can hold only 2
electrons
• Electrons must have opposite spins
(why important?)
• Spin can have two values +½ or –½
Pauli Exclusion Principle
“in a given atom, no two electrons can
have the same set of four quantum
numbers”
What this means for the atom?
• Each atomic sub-orbital may contain a
maximum of 2 electrons
• Those electrons must have opposite
spins
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
6d
5p
4d
4p
5d
3d
3p
2p
Helium with 2
electrons
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
6d
5p
4d
4p
5d
3d
3p
2p
Li with 3 electrons
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
6d
5p
4d
4p
5d
3d
3p
2p
Boron with 5 electrons
5f
4f
2 more important rules:
• Aufbau Principle – electrons enter
orbitals of lowest energy first.
• Hund’s Rule -- When electrons occupy
orbitals of equal energy, one electron
enters each orbital before they pair.
For Example:
2s
2p
After the s sublevel gets
two electrons, three
electrons enter the p
orbitals before they pair.
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
6d
5p
4d
4p
3p
2p
5d
3d
5f
4f
Electron Configuratoin
p
s
d
f
3 QUESTIONS TO ASK
• What Row?
–(principle energy level)
• What section?
–(type of sub-orbital)
• What seat?
–(how many electrons in that
sub-orbital)
Example 1:
Write the electron
configuration for nitrogen.
7N
2
2
3
1s 2s 2p
Example
Write the electron
2:
configuration for Fe.
26Fe
2
2
6
2
6
2
6
Condensed Electron Configurations
• Put the symbol for the Noble gas from
the previous principal energy level, then
add the electron configuration after that
point.
• Example 1 for Nitrogen:
[He] 2s22p3
• Example 2 for Iron:
• [Ar] 4s23d6
Nuclear Chemistry
Chapter 21
Nuclear Radiation
Natural Radioactivity
Nuclear Equations
Producing Radioactive Isotopes
Half-Life
Nuclear Fission and Fusion
Subatomic Particles
• Protons- plus charge
In the nucleus
• Neutrons- neutral
• Electrons - negative charge
Outside the nucleus
Radiation
• Radiation comes from the nucleus of an
atom.
• Unstable nucleus emits a particle or
energy
a alpha
b beta
g gamma
Composition of Stable Nuclei
Notice the nucleus
packs more
neutrons in the
nucleus (compared
to protons) as the
size of the nucleus
increases.
200Hg
80
Number of Neutrons --->
All elements from
#83 and beyond
are naturally
radioactive.
120Sn
50
90Zr
40
20
Number of Protons ----->
83
Radiation Protection
• Shielding
alpha – paper, clothing
beta – lab coat, gloves
gamma- lead, thick concrete
• Limit time exposed
• Keep distance from source
Radiation Protection
Radioactive
particles
and rays
vary greatly
in
penetrating
power.
Alpha Particle a
• Same as a helium nucleus
4
2 He
or a
• Two protons
• Two neutrons
• +2 charge
• Relative penetrating power = 1
Alpha decay
Try this.
• Write the nuclear reaction for alpha
decay of 185Au.
185Au
79
181
Ir +
4
77
Atomic mass decreases by 4,
Atomic No. decreases by 2
because it lost 2 protons and 2
neutrons (a)
a
2
Beta Particle b
An electron emitted from the nucleus
0
e or
b
1
A neutron in the nucleus breaks down
1
1
n
P
0
1
0
+ e
-1
Beta Particle b
• Charge of -1
• Relative Penetrating Power = 100
• Essentially they are fast moving
electrons
Beta decay
234Th
90

234Pa
91
+ 0e
1
beta particle
b
Atomic mass stays the same,
at. no. goes up by 1
because a neutron changes into a proton.
Try this
Write the nuclear equation for the beta
emitter Co 60.
60Co
27
60Ni
28
+
0
e
1
Gamma (g) Radiation
• Pure radiation - high energy wave (photons)
• No mass, NOT shown in the equation
• Almost all radioactive changes involve some
amount of gamma rays being released
• Charge = 0
• Mass = 0
• Relative Penetrating Power = 10,000
Gamma radiation
No change in atomic or mass number
11B
5
11B
5
+
0g
0
boron atom in a
high-energy state
Unnecessary to write a nuclear reaction for this.
Two other common types of
Radioactive Decay
• Positron - same mass as an electron, but
opposite charge
0
e
+1
- has a very short life: annihilated as
soon as it collides with an electron
Write a Nuclear reaction for C-11 positron emission:
11
11
0
C
B +
e
6
5
+1
At. No. decreases by 1 because a proton is converted
to a neutron within the nucleus.
• Electron Capture - an electron from the electron
cloud is captured by the nucleus
0
e
1
Write a nuclear reaction for Rb-81 undergoing electron capture.
81
Rb
37
+
0
e
1
81
Kr
36
At. no. decreases by 1 because the effect is same as
a positron, a proton is converted into a neutron.
Balancing Nuclear Equations
• In the reactants and products
Atomic numbers must balance,
Mass numbers must balance,
and Charges must balance
Half-Life of a Radioisotope
The time for the radiation level to fall
(decay) to one-half its initial value
decay curve
initial
1
half-life
8 mg
2
4 mg
3
2 mg
1 mg
Examples of Half-Life
Isotope
C-15
Ra-224
Ra-223
I-125
C-14
U-235
Half life
2.4 sec
3.6 days
12 days
60 days
5700 years
710 000 000 years
Learning Check
The half life of I-123 is 13 hr.
How much of a 64 mg sample
of I-123 is left after 26 hours?
64/2 = 32 mg remaining in 13 hr.
32/2 = 16 mg remaining after 26 hr.
Answer: 16 mg
Producing Radioactive
Isotopes
Bombardment of atoms produces
radioisotopes
= 60
59Co
+
27
= 60
1n
56Mn
0
25
= 27
cobalt neutron
atom
particle
+
4H e
2
= 27
manganese
radioisotope
alpha
Learning Check
What radioactive isotope is produced in
the following bombardment of boron?
10B
5
+
4He
2
? +
0
1n
Solution
What radioactive isotope is produced in
the following bombardment of boron?
10B
5
+
4He
2
13N
7
nitrogen
radioisotope
+
1n
0
Nuclear Fission
Fission
large nuclei break up
235U
92
+
1n
0
139Ba
56
+
94Kr
36
+ 3 1n Energy
+
0
Fission
Nuclear Fusion
Fusion
small nuclei combine
2H
1
+
3H
4He
1
2
+
1n
+
0
Occurs in the sun and other stars
Energy
Learning Check
Indicate if each of the following are
(1)Fission &/or (2) Fusion
A.
B.
C.
D.
Nucleus splits
Large amounts of energy released
Small nuclei form larger nuclei
Hydrogen nuclei react
Energy
Solution
Indicate if each of the following are
(1)Fission
(2) Fusion
A. 1
Nucleus splits
B. 1 + 2 Large amounts of energy
released
C. 2
Small nuclei form larger nuclei
D. 2
Hydrogen nuclei react