Transcript 1s 2s 2p - Solon City Schools

Atoms, Molecules, and Ions
Chemistry Timeline #1
B.C.
400 B.C. Demokritos and Leucippos use the term "atomos”

2000 years of Alchemy
1500's
 Georg Bauer: systematic metallurgy
 Paracelsus: medicinal application of minerals
1600's
Robert Boyle:The Skeptical Chemist. Quantitative experimentation, identification of
elements
1700s'
 Georg Stahl: Phlogiston Theory
 Joseph Priestly: Discovery of oxygen
 Antoine Lavoisier: The role of oxygen in combustion, law of conservation of
mass, first modern chemistry textbook
Chemistry Timeline #2
1800's
Joseph Proust: The law of definite proportion (composition)
 John Dalton: The Atomic Theory, The law of multiple proportions
Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules
Amadeo Avogadro: Molar volumes of gases
Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements
 Dmitri Mendeleyev: The periodic table
 J.J. Thomson: discovery of the electron
 Henri Becquerel: Discovery of radioactivity
1900's
 Robert Millikan: Charge and mass of the electron
 Ernest Rutherford: Existence of the nucleus, and its relative size
 Meitner & Fermi: Sustained nuclear fission
 Ernest Lawrence: The cyclotron and trans-uranium elements
Laws
• Conservation of Mass
• Law of Definite Proportion –
– compounds have a constant composition.
– They react in specific ratios by mass.
• Multiple Proportions-
– When two elements form more than one
compound, the ratios of the masses of the
second element that combine with one
gram of the first can be reduced to small
whole numbers.
Proof
• Mercury has two oxides.
– One is 96.2 % mercury by mass, the
other is 92.6 % mercury by mass.
• Show that these compounds follow
the law of multiple proportion.
• Speculate on the formula of the two
oxides.
Dalton’s Atomic Theory (1808)
 All matter is composed of extremely
small particles called atoms
 Atoms of a given element are
identical in size, mass, and other
properties; atoms of different
John Dalton
elements differ in size, mass, and
other properties
 Atoms cannot be subdivided, created, or destroyed
 Atoms of different elements combine in simple
whole-number ratios to form chemical compounds
 In chemical reactions, atoms are combined,
separated, or rearranged
Modern Atomic Theory
Several changes have been made to Dalton’s theory.
Dalton said:
Atoms of a given element are identical in
size, mass, and other properties; atoms of
different elements differ in size, mass, and
other properties
Modern theory states:
Atoms of an element have a characteristic
average mass which is unique to that
element.
Modern Atomic Theory #2
Dalton said:
Atoms cannot be subdivided, created, or destroyed
Modern theory states:
Atoms cannot be subdivided, created, or destroyed
in ordinary chemical reactions. However, these
changes CAN occur in nuclear reactions!
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube
to deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
Thomson’s Atomic Model
Thomson believed that the electrons were like plums
embedded in a positively charged “pudding,” thus it was
called the “plum pudding” model.
Rutherford’s Gold Foil Experiment
 Alpha particles are helium nuclei
 Particles were fired at a thin sheet of gold foil
 Particle hits on the detecting screen (film) are
recorded
Atomic Particles
Particle
Charge
Mass (kg)
Location
Electron
-1
9.109 x 10-31
Electron
cloud
Proton
+1
1.673 x 10-27
Nucleus
0
1.675 x 10-27
Nucleus
Neutron
The Atomic
Scale
 Most of the mass of the
atom is in the nucleus
(protons and neutrons)
 Electrons are found
outside of the nucleus (the
electron cloud)
 Most of the volume of
the atom is empty space
“q” is a particle called a “quark”
About Quarks…
Protons and neutrons are
NOT fundamental particles.
Protons are made of
two “up” quarks and
one “down” quark.
Neutrons are made of
one “up” quark and
two “down” quarks.
Quarks are held together
by “gluons”
Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Nucleus
Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Carbon = 12.011
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.11%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
Isotope
Molecules
Two or more atoms of the same or different
elements, covalently bonded together.
Molecules are discrete structures, and their
formulas represent each atom present in the
molecule.
Benzene, C6H6
Covalent Network Substances
Covalent network substances have covalently
bonded atoms, but do not have discrete
formulas.
Why Not??
Graphite
Diamond
Ions
 Cation: A positive ion
• Mg2+, NH4+
 Anion: A negative ion
 Cl-, SO42-
 Ionic Bonding: Force of attraction between
oppositely charged ions.
 Ionic compounds form crystals, so their
formulas are written empirically (lowest whole
number ratio of ions).
Periodic Table with Group Names
Predicting Ionic Charges
Group 1: Lose 1 electron to form 1+ ions
H+
Li+ Na+
K+
Predicting Ionic Charges
Group 2: Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+
Predicting Ionic Charges
B3+
Al3+
Ga3+
Group 13: Loses 3
electrons to form
3+ ions
Predicting Ionic Charges
Caution! C22- and C4are both called carbide
Group 14: Loses 4
electrons or gains
4 electrons
Predicting Ionic Charges
N3- Nitride
P3- Phosphide
As3- Arsenide
Group 15: Gains 3
electrons to form
3- ions
Predicting Ionic Charges
O2- Oxide
S2- Sulfide
Se2- Selenide
Group 16: Gains 2
electrons to form
2- ions
Predicting Ionic Charges
F- Fluoride
Br- Bromide
Cl- Chloride
I- Iodide
Group 17: Gains 1
electron to form
1- ions
Predicting Ionic Charges
Group 18: Stable
Noble gases do not
form ions!
Predicting Ionic Charges
Groups 3 - 12: Many transition elements
have more than one possible oxidation state.
Iron(II) = Fe2+
Iron(III) = Fe3+
Predicting Ionic Charges
Groups 3 - 12: Some transition elements
have only one possible oxidation state.
Zinc = Zn2+
Silver = Ag+
Cadmium = Cd2+
Writing Ionic Compound Formulas
Example: Barium nitrate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
2+
Ba ( NO3- ) 2
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
Not balanced!
Writing Ionic Compound Formulas
Example: Ammonium sulfate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
( NH4+) SO42-
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
2
Not balanced!
Naming Ionic Compounds
• 1. Cation first, then anion
• 2. Monatomic cation = name of the element
• Ca2+ = calcium ion
• 3. Monatomic anion = root + -ide
• Cl- = chloride
• CaCl2 = calcium chloride
Naming Ionic Compounds
(continued)
Metals with multiple oxidation states
 some metal forms more than one cation
 use Roman numeral in name
PbCl2
Pb2+ is the lead(II) cation
PbCl2 = lead(II) chloride
Naming Binary Compounds





Compounds between two nonmetals
First element in the formula is named first.
Second element is named as if it were an anion.
Use prefixes
Only use mono on second element P2O5 = diphosphorus pentoxide
CO2 = carbon dioxide
CO = carbon monoxide
N2O = dinitrogen monoxide
Acids
• Substances that produce H+ ions
when dissolved in water.
• All acids begin with H.
• Two types of acids:
• Oxyacids
• Non-oxyacids
Naming acids
• If the formula has oxygen in it
• write the name of the anion, but
change
– ate to -ic acid
– ite to -ous acid
• Watch out for sulfuric and sulfurous
• H2CrO4
• HMnO4
• HNO2
Naming acids
•
•
•
•
•
•
If the acid doesn’t have oxygen
add the prefix hydrochange the suffix -ide to -ic acid
HCl
H2S
HCN
Formulas for acids
•
•
•
•
•
•
•
Hydrofluoric acid
Dichromic acid
Carbonic acid
Hydrophosphoric acid
Nitric acid
Perchloric acid
Phosphorous acid
HF
H2Cr2O7
H2CO3
H3P
HNO3
HClO4
H3PO3
Selenium would commonly form
this ion:
Se2+
Se+
Se2Sl2S2SeSe36
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Se
3-
5
0%
Se
-
4
0%
S2
-
3
Se
+
2
Se
2+
1
0%
Sl2
-
0%
Se
2-
1.
2.
3.
4.
5.
6.
7.
Cesium would commonly form
this ion:
Ce2+
Cs+
Cs2CCsCs2+
Cm+
7
8
9
10
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0%
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0%
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+
6
0%
Cm
5
0%
Cs
2+
4
0%
Cs
-
3
Cs
+
2
Ce
2+
1
0%
C-
0%
Cs
2-
1.
2.
3.
4.
5.
6.
7.
This is the formula for zinc
hydroxide:
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10
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H)
2
18
19
0%
Zn
2H
8
0%
Zn
2(
O
7
OH
2)
2
6
0%
Zn
(
5
0%
Zn
H2
4
0%
OH
)2
3
Zn
O
2
Zn
O
1
0%
H2
0%
Zn
(
ZnOH
ZnOH2
Zn(OH)2
ZnH2
Zn(OH2)2
Zn2(OH)2
Zn2H
H
1.
2.
3.
4.
5.
6.
7.
20
This is the formula for
hydrochloric acid:
8
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10
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30
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12
13
14
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18
0%
H2
Cl
O
7
0%
H2
Cl
6
4
5
0%
HC
lO
4
3
3
0%
HC
lO
2
0%
2
1
0%
HC
lO
0%
HC
lO
HCl
HClO
HClO2
HClO3
HClO4
H2Cl
H2ClO
HC
l
1.
2.
3.
4.
5.
6.
7.
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20
Iron would commonly form this
ion:
Fe2+
Fe+
Fe2FeIr2+
Ir+
Fe3+
6
7
8
9
10
21
22
23
24
25
26
27
28
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30
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12
13
0%
Fe
3+
5
0%
Ir+
4
0%
Ir2
+
3
0%
Fe
-
2
0%
Fe
2-
1
0%
Fe
+
0%
Fe
2+
1.
2.
3.
4.
5.
6.
7.
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This slide contains
classified material and
cannot be shown to high
school students. Please
continue as if everything is
normal.
Which points of Dalton’s theory are not true based
on current understanding of the atom?
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...
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m
m
so
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at
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3.
At
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2.
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0%
1.
Al
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0%
sc
an
1. All matter is composed of extremely small
particles called atoms
2. Atoms of a given element are identical in size,
mass, and other properties; atoms of different
elements differ in size, mass, and other
properties
3. Atoms cannot be subdivided, created, or
destroyed
4. Atoms of different elements combine in simple
whole-number ratios to form chemical
compounds
5. In chemical reactions, atoms are combined,
separated, or rearranged
19
20
Quantum Mechanics
The Puzzle of the Atom
 Protons and electrons are attracted to each
other because of opposite charges
 Electrically charged particles moving in a
curved path give off energy
 Despite these facts, atoms don’t collapse
Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the
electron as a particle.
His son, George Thomson won the Nobel prize for
describing the wave-like nature of the electron.
The
electron is
a particle!
The
electron is
an energy
wave!
Toupee?
The Wave-like Electron
The electron propagates
through space as an energy
wave. To understand the
atom, one must understand
the behavior of
electromagnetic waves.
Louis deBroglie
Electromagnetic radiation propagates
through space as a wave moving at the
speed of light.
c = 
C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters
Types of electromagnetic radiation:
The energy (E ) of electromagnetic
radiation is directly proportional to the
frequency () of the radiation.
E = h
E = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x 10-34 J·s)
 = frequency, in units of hertz (hz, sec-1)
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High Frequency
=
High ENERGY
Wavelength Table
Relating Frequency, Wavelength
and Energy
c 
E  h
Common re-arrangements:
E
hc

hc

E
Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum
Spectroscopic analysis of the hydrogen
spectrum…
…produces a “bright line” spectrum
Electron transitions
involve jumps of
definite amounts of
energy.
This produces bands
of light with definite
wavelengths.
Bohr Model Energy Levels
Electron Energy in Hydrogen
Eelectron  - 2.178 x 10
-18
Z
J  2
n
2



Z = nuclear charge (atomic number)
n = energy level
***Equation works only for atoms or ions
with 1 electron (H, He+, Li2+, etc).
Calculating Energy Change, E, for
Electron Transitions
2
2 

Z 
-18  Z
E  - 2.178 x 10 J 2 - 2
n

n
initial 
 final
Energy must be absorbed from a photon
(+E) to move an electron away from the
nucleus
Energy (a photon) must be given off (-E)
when an electron moves toward the nucleus
Quantum Numbers
Each electron in an atom has a unique
set of 4 quantum numbers which describe
it.
 Principal quantum number
(n)
 Angular momentum quantum number (l)
 Magnetic quantum number (m)
 Spin quantum number
(s)
Pauli Exclusion Principle
No two electrons in an atom
can have the same four
quantum numbers.
Wolfgang
Pauli
Principal Quantum Number
Generally symbolized by n, it denotes the shell
(energy level) in which the electron is located.
Number of electrons
that can fit in a shell:
2n2
Angular Momentum Quantum Number
The angular momentum quantum number, generally
symbolized by l, denotes the orbital (subshell) in
which the electron is located.
l =3
f
Magnetic Quantum Number
The magnetic quantum number, generally
symbolized by m, denotes the orientation of the
electron’s orbital with respect to the three axes in
space.
Assigning the Numbers
 The three quantum numbers (n, l, and m) are
integers.
 The principal quantum number (n) cannot be
zero.
 n must be 1, 2, 3, etc.
 The angular momentum quantum number (l )
can be any integer between 0 and n - 1.
 For n = 3, l can be either 0, 1, or 2.
 The magnetic quantum number (ml) can be any
integer between -l and +l.
 For l = 2, m can be either -2, -1, 0, +1, +2.
Principle, angular momentum, and magnetic
quantum numbers: n, l, and ml
Spin Quantum Number
Spin quantum number denotes the behavior
(direction of spin) of an electron within a
magnetic field.
Possibilities for electron spin:
1

2
1
2
An orbital is a region within an atom where there
is a probability of finding an electron. This is a
probability diagram for the s orbital in the first
energy level…
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.
Schrodinger Wave Equation

d
V 
8  m dx
h
2
2
2
2
 E
Equation for probability of a
single electron being found
along a single axis (x-axis)
Erwin Schrodinger
Heisenberg Uncertainty Principle
“One cannot simultaneously
determine both the position
and momentum of an electron.”
You can find out where the
electron is, but not where it
is going.
Werner
Heisenberg
OR…
You can find out where the
electron is going, but not
where it is!
Sizes of s orbitals
Orbitals of the same shape (s, for instance) grow
larger as n increases…
Nodes are regions of low probability within an
orbital.
Orbitals in outer energy levels DO penetrate into
lower energy levels. Penetration #1
This is a probability
Distribution for a
3s orbital.
What parts of the
diagram correspond
to “nodes” – regions
of zero probability?
Which of the orbital types in the 3rd energy level
Does not seem to have a “node”?
WHY NOT?
Penetration #2
The s orbital has a spherical shape centered around
the origin of the three axes in space.
s orbital shape
P orbital shape
There are three peanut-shaped p orbitals in
each energy level above n = 1, each assigned to
its own axis (x, y and z) in space.
d orbital shapes
Things get a bit more
complicated with the five d
orbitals that are found in
the d sublevels beginning
with n = 3. To remember
the shapes, think of:
“double peanut”
…and a “peanut
with a donut”!
Shape of f orbitals
Things get even more
complicated with the seven f
orbitals that are found in the
f sublevels beginning with n
= 4. To remember the
shapes, think of:
Flower
Element
Lithium
Configuration
notation
1s22s1
[He]2s1
____
1s
Beryllium
____
____
2p
____
____
2s
____
____
2p
____
[He]2s2p2
____
2s
____
____
2p
____
1s22s2p3
[He]2s2p3
____
2s
____
____
2p
____
1s22s2p4
[He]2s2p4
____
2s
____
____
2p
____
1s22s2p5
[He]2s2p5
____
1s
Neon
____
2s
1s22s2p2
____
1s
Fluorine
____
[He]2s2p1
____
1s
Oxygen
____
2p
1s22s2p1
____
1s
Nitrogen
____
[He]2s2
____
1s
Carbon
____
2s
1s22s2
____
1s
Boron
Noble gas
notation
Orbital notation
____
2s
____
____
2p
____
1s22s2p6
[He]2s2p6
____
1s
____
2s
____
____
2p
____
Orbital filling table
Electron configuration of the
elements of the first three series
Irregular confirmations of Cr and Cu
Chromium steals a 4s electron to half
fill its 3d sublevel
Copper steals a 4s electron to FILL
its 3d sublevel
In Bohr’s atomic theory, when an electron
moves from one energy level to another
energy level more distant from the nucleus.
1.
2.
3.
4.
5.
energy is emitted
energy is absorbed
no change in energy occurs
light is emitted
none of these
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ne
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Which form of electromagnetic
radiation has the longest wavelengths?
io
tio
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gamma rays
microwaves
radio waves
infrared radiation
x-rays
ga
m
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How many electrons in an atom can
have the quantum numbers n = 3, l = 2?
1
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7
8
9
10
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2
1.
2.
3.
4.
5.
20
Which of the following combinations
of quantum numbers is not allowed?
½
–½
½
–½
½
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0½
0
0
–1
–2
0
43
–2
–½
1
0
1
3
2
21
–1
½
1
3
2
4
4
30
0–
½
s
11
0½
m
s
l
nlm
1.
2.
3.
4.
5.
n
20
1.
2.
3.
4.
5.
The electron configuration of
indium is
1s22s22p63s23p64s23d104p65s24d105p15d10
1s22s22p63s23p64s23d104d104p1
1s23s22p63s23p64s24d104p65s25d105p1
1s22s22p63s23p64s23d104p65s24d105p1
none of these
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
13
14
15
0%
0%
0%
0%
1s
22
s2
2p
63
s2
1s
3p
22
64
s2
s2
2p
3d
63
1.
s2
1s
..
3p
23
64
s2
s2
2p
3d
63
1.
s2
1s
..
3p
22
6
s2
4s
2p
24
63
d1
s2
...
3p
64
s2
3d
1.
..
no
ne
of
th
es
e
0%
16
17
18
19
20
Ag has __ electrons in its d orbitals.
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
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13
14
15
16
17
18
19
20
Periodicity
Determination of Atomic Radius:
Half of the distance between nuclei in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period
Increased effective nuclear charge due
to decreased shielding
Radius increases down a group
Addition of principal quantum levels
Table of
Atomic
Radii
Ionization Energy - the energy required to
remove an electron from an atom
 Increases for successive electrons taken from
the same atom
 Tends to increase across a period
Electrons in the same quantum level do
not shield as effectively as electrons in
inner levels
Irregularities at half filled and filled
sublevels due to extra repulsion of
electrons paired in orbitals, making them
easier to remove
 Tends to decrease down a group
Outer electrons are farther from the
nucleus
Ionization of Magnesium
Mg + 738 kJ  Mg+ + eMg+ + 1451 kJ  Mg2+ + eMg2+ + 7733 kJ  Mg3+ + e-
Table of 1st
Ionization Energies
Another Way to Look at Ionization Energy
Yet Another Way to Look at Ionization Energ
Electron Affinity - the energy change
associated with the addition of an electron
 Affinity tends to increase across a period
 Affinity tends to decrease as you go down
in a period
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals
Table of Electron Affinities
Summary of Periodic Trends