Atoms, Molecules, and Ions

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Atoms, Molecules, and Ions
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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 BoyleThe 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

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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

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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.

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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.

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Daltons 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
  elements differ in size, mass, and other
  properties

John Dalton

• 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

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Modern Atomic Theory
Several changes have been made to Daltons 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.
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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!
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Atomic Particles
Particle Charge Mass (kg) Location
Electron -1 9.109 x 10-31 Electron cloud
Proton 1 1.673 x 10-27 Nucleus
Neutron 0 1.675 x 10-27 Nucleus
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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
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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
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Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of
neutrons.
Isotope Protons Electrons Neutrons Nucleus
Hydrogen1 (protium) 1 1 0
Hydrogen-2 (deuterium) 1 1 1
Hydrogen-3 (tritium) 1 1 2
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Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Carbon 12.011
Isotope 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 lt0.01
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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
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Covalent Network Substances
Covalent network substances have covalently
bonded atoms, but do not have discrete formulas.
Why Not??
Graphite
Diamond
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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).

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Periodic Table with Group Names
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cannot be shown to high school students. Please
continue as if everything is normal.
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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.
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Thomsons Atomic Model
Thomson believed that the electrons were like
plums embedded in a positively charged pudding,
thus it was called the plum pudding model.
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Rutherfords 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

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Quantum Mechanics
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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 dont collapse

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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
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Types of electromagnetic radiation
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Long Wavelength Low Frequency Low ENERGY
Wavelength Table
Short Wavelength High Frequency High
ENERGY
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The Wave-like Electron
The electron propagates through space on an
energy wave. To understand the atom, one must
understand the behavior of electromagnetic waves.
Louis deBroglie
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The Great Niels Bohr (1885 - 1962)
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Spectroscopic analysis of the visible spectrum
produces all of the colors in a continuous
spectrum
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Spectroscopic analysis of the hydrogen spectrum
produces a bright line spectrum
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Electron transitionsinvolve jumps of definite
amounts ofenergy.
This produces bands of light with definite
wavelengths.
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Bohr Model Energy Levels
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Schrodinger Wave Equation
Equation for probability of a single electron
being found along a single axis (x-axis)
Erwin Schrodinger
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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.
OR
You can find out where the electron is going, but
not where it is!
Werner Heisenberg
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Quantum Numbers
Each electron in an atom has a unique set of 4
quantum numbers which describe it.

• Principal quantum number
• Angular momentum quantum number
• Magnetic quantum number
• Spin quantum number

(n)
(l)
(m)
(s)
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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
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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
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Magnetic Quantum Number
The magnetic quantum number, generally symbolized
by m, denotes the orientation of the electrons
orbital with respect to the three axes in space.
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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.

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Principle, angular momentum, and magnetic quantum
numbers n, l, and ml
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Pauli Exclusion Principle
No two electrons in an atom can have the same
four quantum numbers.
Wolfgang Pauli
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Spin Quantum Number
Spin quantum number denotes the behavior
(direction of spin) of an electron within a
magnetic field.
Possibilities for electron spin
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An orbital is a region within an atom where
thereis 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.
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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.
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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?
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The s orbital has a spherical shape centered
around the origin of the three axes in space.
s orbital shape
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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.
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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!
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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
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Orbital filling table
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Element Configuration notation Orbital notation Noble gas notation
Lithium 1s22s1 ____ ____ ____ ____ ____ 1s 2s 2p He2s1
Beryllium 1s22s2 ____ ____ ____ ____ ____ 1s 2s 2p He2s2
Boron 1s22s2p1 ____ ____ ____ ____ ____ 1s 2s 2p He2s2p1
Carbon 1s22s2p2 ____ ____ ____ ____ ____ 1s 2s 2p He2s2p2
Nitrogen 1s22s2p3 ____ ____ ____ ____ ____ 1s 2s 2p He2s2p3
Oxygen 1s22s2p4 ____ ____ ____ ____ ____ 1s 2s 2p He2s2p4
Fluorine 1s22s2p5 ____ ____ ____ ____ ____ 1s 2s 2p He2s2p5
Neon 1s22s2p6 ____ ____ ____ ____ ____ 1s 2s 2p He2s2p6
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Electron configuration of the elements of the
first three series
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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
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In Bohrs atomic theory, when an electron moves
from one energy level to another energy level
more distant from the nucleus.

1. energy is emitted
2. energy is absorbed
3. no change in energy occurs
4. light is emitted
5. none of these

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
21 22 23 24 25 26 27 28 29 30 31 32
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Which form of electromagnetic radiation has the
longest wavelengths?

1. gamma rays
2. microwaves
3. radio waves
4. infrared radiation
5. x-rays

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
21 22 23 24 25 26 27 28 29 30 31 32
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How many electrons in an atom can have the
quantum numbers n 3, l 2?

1. 2
2. 5
3. 10
4. 18
5. 6

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
21 22 23 24 25 26 27 28 29 30 31 32
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Which of the following combinations of quantum
numbers is not allowed?

• n l m s
• 1 1 0 ½
• 3 0 0 ½
• 2 1 1 ½
• 4 3 2 ½
• 4 2 0 ½

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21 22 23 24 25 26 27 28 29 30 31 32
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The electron configuration of indium is

1. 1s22s22p63s23p64s23d104p65s24d105p15d10
2. 1s22s22p63s23p64s23d104d104p1
3. 1s23s22p63s23p64s24d104p65s25d105p1
4. 1s22s22p63s23p64s23d104p65s24d105p1
5. none of these

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
21 22 23 24 25 26 27 28 29 30 31 32
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Ag has __ electrons in its d orbitals.





1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
21 22 23 24 25 26 27 28 29 30 31 32
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Periodicity
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Atomic Size

Radius

• Atomic Radius half the distance between two
  nuclei of a diatomic molecule.

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Trends in Atomic Size

• Influenced by three factors.
• Energy Level
• Higher energy level is further away.
• Charge on nucleus
• More charge pulls electrons in closer.
• Shielding
• Layers of electrons shield from nuclear pull.

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Shielding

• The electron on the outside energy level has to
  look through all the other energy levels to see
  the nucleus

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Shielding

• The electron on the outside energy level has to
  look through all the other energy levels to see
  the nucleus.
• A second electron has the same shielding.

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Group trends
H

• As we go down a group
• Each atom has another energy level,
• So the atoms get bigger.

Li
Na
K
Rb
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Periodic Trends

• As you go across a period the radius gets
  smaller.
• Same energy level.
• More nuclear charge.
• Outermost electrons are closer.

Na
Mg
Al
Si
P
S
Cl
Ar
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Table of Atomic Radii
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Ionic Size

• Cations form by losing electrons.
• Cations are smaller that the atom they come from.
• Metals form cations.
• Cations of representative elements have noble gas
  configuration.

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Ionic size

• Anions form by gaining electrons.
• Anions are bigger that the atom they come from.
• Nonmetals form anions.
• Anions of representative elements have noble gas
  configuration.

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Overall
K
Na
Li
Atomic Radius (nm)
Kr
Ar
Ne
H
10
Atomic Number
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Ionization Energy

• The amount of energy required to completely
  remove an electron from a gaseous atom.
• Removing one electron makes a 1 ion.
• The energy required is called the first
  ionization energy.

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Ionization Energy

• The second ionization energy is the energy
  required to remove the second electron.
• Always greater than first IE.
• The third IE is the energy required to remove a
  third electron.
• Greater than 1st of 2nd IE.

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Symbol First Second Third
11810 14840 3569 4619 4577
5301 6045 6276
1312 2731 520 900 800 1086 1402 1314 1681 2080
H HeLi BeB C N O F Ne
5247 7297 1757 2430 2352 2857 3391 3375 3963
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What determines IE

• The greater the nuclear charge the greater IE.
• Distance from nucleus increases IE
• Filled and half filled orbitals have lower
  energy, so achieving them is easier, lower IE.
• Shielding

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Group trends

• As you go down a group first IE decreases because
• The electron is further away.
• More shielding.

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Periodic trends

• All the atoms in the same period have the same
  energy level.
• Same shielding.
• Increasing nuclear charge
• So IE generally increases from left to right.
• Exceptions at full and 1/2 fill orbitals.

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He

• He has a greater IE than H.
• same shielding
• greater nuclear charge

H
First Ionization energy
Atomic number
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He

• Li has lower IE than H
• more shielding
• further away
• outweighs greater nuclear charge

H
First Ionization energy
Li
Atomic number
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He

• Be has higher IE than Li
• same shielding
• greater nuclear charge

H
First Ionization energy
Be
Li
Atomic number
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He

• B has lower IE than Be
• same shielding
• greater nuclear charge
• By removing an electron we make s orbital half
  filled

H
First Ionization energy
Be
B
Li
Atomic number
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He
C
H
First Ionization energy
Be
B
Li
Atomic number
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He
N
C
H
First Ionization energy
Be
B
Li
Atomic number
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He

• Breaks the pattern because removing an electron
  gets to 1/2 filled p orbital

N
O
C
H
First Ionization energy
Be
B
Li
Atomic number
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He
F
N
O
C
H
First Ionization energy
Be
B
Li
Atomic number
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He
Ne

• Ne has a lower IE than He
• Both are full,
• Ne has more shielding
• Greater distance

F
N
O
C
H
First Ionization energy
Be
B
Li
Atomic number
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He
Ne

• Na has a lower IE than Li
• Both are s1
• Na has more shielding
• Greater distance

F
N
O
C
H
First Ionization energy
Be
B
Li
Na
Atomic number
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First Ionization energy
Atomic number
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Driving Force

• Full Energy Levels are very low energy.
• Noble Gases have full orbitals.
• Atoms behave in ways to achieve noble gas
  configuration.

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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
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Table of Electron Affinities
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Electronegativity
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Electronegativity

• The tendency for an atom to attract electrons to
  itself when it is chemically combined with
  another element.
• How fair it shares.
• Big electronegativity means it pulls the electron
  toward it.
• Atoms with large negative electron affinity have
  larger electronegativity.

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Group Trend

• The further down a group the farther the electron
  is away and the more electrons an atom has.
• More willing to share.
• Low electronegativity.

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Periodic Trend

• Metals are at the left end.
• They let their electrons go easily
• Low electronegativity
• At the right end are the nonmetals.
• They want more electrons.
• Try to take them away.
• High electronegativity.

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Ionization energy, electronegativity Electron
affinity INCREASE
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Atomic size increases, shielding constant
Ionic size increases
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Another Way to Look at Ionization Energy
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Yet Another Way to Look at Ionization Energy
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Summary of Periodic Trends