Download Chapter 2 Atoms and Elements Modern Atomic theory

Principles of Chemistry:
A Molecular Approach,
1st Ed.
Nivaldo Tro
Chapter 2
Atoms and Elements
Chapt 2
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Modern Atomic theory and the Laws that
led to it
Atomic Theory- “All matter is composed of atoms
grew out of observation and laws”
Three important Laws that led to development and
acceptance of Atomic Theory
Law of conservation of mass
Law of definite proportions
Law of multiple proportions
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Law of Conservation of Mass
In a chemical reaction, matter is neither created nor
destroyed. (Antoine Lavoisier1743–1794)
Total mass of the materials you have before the
reaction must equal the total mass of the materials
you have at the end.
total mass of reactants = total mass of products
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1
Reaction of Sodium with Chlorine to Make
Sodium Chloride
The mass of sodium and chlorine used is determined by the number of
atoms that combine
7.7 g Na
+ 11.9 g Cl2
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19.6 g NaCl
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Law of Definite Proportions
All samples of a given compound, regardless of their
source or how they were prepared, have the same
proportions of their constituent elements. Joseph
Proust (1754–1826)
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Show that two samples of carbon dioxide
obey the Law of Definite Proportions
Given: Sample 1: 25.6 g O and 9.60 g C
Sample 2: 21.6 g O and 8.10 g C
Find: Proportion O:C
Conceptual
g O1 , g C 1
O:C in each sample
Plan:
g O2 , g C 2
Relationships:
All samples of a compound have the same
proportion of elements by mass.
Solution:
Check: Since both samples have the same O:C ratio, the result
is consistent with the Law of Definite Proportions.
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Law of Definite Proportions
Practice—If a 10.0-g sample of calcite contains 4.0 g
of calcium, how much calcite contains 0.24 g of
calcium?
Given:Sample 1: 4.0 g Ca and 10.0 g calcite
Sample 2: 0.24 g Ca
Find :mass calcite, g
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Law of Multiple Proportions
When two elements A and B form two different compounds, the
masses of B that combine with 1 g of A can be expressed as a
ratio of small whole numbers. (John Dalton1766–1844)
For example oxides of carbon (CO and CO2)
CO
CO2
Ratio of these two masses itself is a small whole
number
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Dalton’s Atomic Theory
1)
2)
3)
4)
Dalton explained these laws with his atomic theory which
included the following concepts;
Each element is composed of tiny, indestructible particles called
atoms.
All atoms of a given element have the same mass and other
properties that distinguish them from atoms of other elements
Atoms combine in simple, whole-number ratios to form
compounds
Atoms of one element cannot change into atoms of another
element.,
In a chemical reaction atoms simply rearrange the way they
are attached
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The Discovery of electron
Although Dalton postulated that atoms were indivisible,
experiments at the beginning of this century showed
that atom itself was composed of even smaller, more
fundamental particles.
J.J. Thompson In 1897 British physicist
Cathode ray tube identified the
electrons(subatomic particles) See Figure 2.3
calculated the ratio of the electron’s mass, m e, to
its electric charge, e (m/e)
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Cathode Ray Tube
•An evacuated glass tube containing metal electrodes
when connected to a high-voltage power supply,
produced a stream of particles called cathode rays.
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Properties of cathode rays
Cathode rays- are emitted from negatively charged
electrode, called cathode travel to positively charged
electrode called anode
Cathode rays are
stream of negatively charged particles
They traveled in straight lines
were independent of the composition of the
material from which they originated(cathode)
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Some Notes on Charge
opposite charges attract
like charges repel
To be neutral, something must have
no charge or equal amounts of opposite charges.
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Thomson’s Results
The cathode rays are stream of negatively charged
tiny particles because the beam always deflected
toward the + plate
The amount of deflection was related to two factors;
the charge and mass of the particles
Every material tested contained these same particles
The charge:mass ratio of these particles was -1.76
× 108 C/g
2000 times lighter than hydrogen the lightest atom known
Millikan showed that the particle did have the
same amount of charge as the hydrogen ion
became known as electrons
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Millikan’s Oil Drop Experiment
In 1909, U.S. physicist, Robert Millikan had obtained
the charge on the electron
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Millikan’s Oil Drop Experiment
Oil droplets received charge from ionizing radiation
Charged drops suspended in an electric field
The mass and charge of each droplet were used to calculate the
mass and charge of single electron
Electrons
Electrons are particles found in all atoms.
Cathode rays are streams of electrons.
The electron has a charge of -1.60 × 1019 C.
The electron has a mass of 9.1 × 10-28 g.
Charge x mass /charge = mass
-1.60 × 1019 C x g/-1.76 x 10 8C = 9.1 × 10-28 g.
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A New Theory of the Atom
Thomson’s new model of the atom replaced the first
statement in Dalton’s Atomic Theory. He proposed
that atom actually has an inner structure
atom is no longer indivisible
Plum-Pudding Modle (J.J. Thompson)
Nuclear Atom (Rutherford’s Gold foil expt.)
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Thomson’s Plum Pudding Atom
The structure of the atom contains many negatively
charged electrons held within a positively charged
sphere (atom)
There had to be a source of positive charge because the
atom is neutral
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Radioactivity
Henri Becquerel and Marie Curie(1800), certain
elements would constantly emit small, energetic
particles and rays.
These energetic particles could penetrate matter.
three different kinds of emissions.
alpha, , particles with a mass 4× H atom and +
charge
beta, , particles with a mass ~1/2000 H atom
and – charge
gamma, , rays that are energy rays, not particles
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Structure of the atom
In 1897, the British physicist J.J. Thompson showed that
atoms were not indivisible particles.
Thompson calculated the ratio of the electron’s mass, m e, to
its electric charge, e.
In 1909, U.S. physicist, Robert Millikan had obtained the
charge on the electron.
These two discoveries combined provided us with the
electron’s mass of 9.109 x 10-31 kg, which is more than
2000 times smaller than the mass of the lightest atom
(hydrogen).
These experiments showed that the electron was indeed a
subatomic particle
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Rutherford’s Experiment
Idea of the nuclear model of the atom based on
experiments done in his laboratory by Hans Geiger
and Ernest Morrison. (See Figure 2.5)
Rutherford’s famous gold Foil experiment
majority of the atom’s mass and all of its
positive charge are contained in a small core
called nucleus.
Alpha particles directed at a thin gold film deflected
in all directions including back at the source
Only a large positive charge could cause the –
particles to bounce back
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Rutherford’s Nuclear Atom
An atom is mostly empty space throughout which tiny
negatively charged electrons are dispersed
because almost all the particles went straight through
Most of atoms ‘s mass and all of its positive charge are
contained in a small core called nucleus.
because of the few particles that bounced back
This dense particle was positively charged.
There are as many negatively charged electrons outside the
nucleus as there are positively charged particles(protons)
within the the nucleus, so the atom is electrically neutral.
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Plum Pudding
Atom
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
A few of the
•
•
particles
do not go through. Nuclear Atom
.
Almost all particles
go straight through.
.
.
If an atom was like
a plum pudding,
all the particles
should go
straight through.
Some particles(1 in 20 000)
go through, but are deflected due to
+:+ repulsion from the nucleus.
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Structure of the Atom
Protons-has charge same as an electron but with
opposite sign.
Since protons and electrons have the same amount
of charge, for the atom to be neutral there, must
be equal numbers of protons and electrons.
The nucleus of an atom is composed of two different
kinds of particles, protons and neutrons.
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Subatomic particles
Neutron ( Rutherford and Chadwick)
have no charge
a mass of 1 amu,slightly heavier than a proton
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Elements
Atomic number (Z)-The number of protons in the
nucleus of an atom
Each element has a unique number of protons (Z)
The elements are arranged on the Periodic Table
in order of their atomic numbers.
Each element has a unique name and symbol.
symbol either one or two letters of the name in
english , Greek or Latin
one capital letter or one capital letter and one
lowercase
e.g. Hydrogen = H for sodium(natrium Latin)= Na
Tin = Sn (stannum in Latin)
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Atomic number
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Isotopes
Isotopes- Atoms of the same element with same
number of protons but different number of neutrons
e.g for all carbon atoms Z= 6 (C-12 and C-13)
The observed mass is a weighted average of the
weights of all the naturally occurring atoms
Natural abundance- The percentage of one
isotope in an element
The atomic mass of chlorine is 35.45 amu
35
17 Cl
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17 Cl
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Isotopes
All isotopes of an element are chemically identical.
undergo the exact same chemical reactions
have the same number of protons.
have different masses due to different numbers of
neutrons.
Isotopes are identified by their mass numbers (A) .
Mass number A = protons (p) + neutrons (n)
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Isotopes
• A nuclide is an atom characterized by a definite
atomic number and mass number.
The shorthand notation for a sodium nuclide
• Atomic Number Z
=Number of protons
Mass number= neutron + proton
• Mass Number A
= Protons + Neutrons
sodium
23
11 Na
23
Atomic number = # proton
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Neon
Symbol
Number of Number of A, Mass
Protons
Neutrons Number
Percent
Natural
Abundance
Ne-20 or 20
10 Ne
10
10
20
90.48%
21Ne
Ne-21 or 10
10
11
21
0.27%
22 Ne
Ne-22 or 10
10
12
22
9.25%
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Practice—Complete the Table
Atomic
Number
Protons
Neutrons
6
7
Mass
Number
Atomic
Symbol
Electrons
42
96
55
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Ions: losing and gaining electrons
Ions- When atoms gain or lose electrons, they
acquire a charge and are called ions
Anions –Negatively charged ions
(atoms gain electrons)
Cations - Positively charged ions
(atoms lose electrons)
Ions behave much differently than the neutral atom.
Since materials like table salt are neutral, there
must be equal amounts of charge from cations
and anions in them.
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Atomic Structures of Ions
Nonmetals form anions
For each negative charge, the ion has 1 more
electron than the neutral atom.
F = 9 p+ and 9 e-, F- = 9 p+ and 10 eAnions are named by changing the ending of the
name to -ide
fluorine
F + 1eFfluoride ion
oxygen
O + 2eO2oxide ion
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IONS
Metals form cations
For each positive charge, the ion has 1 less electron
than the neutral atom.
Na atom = 11 p+ and 11 e-, Na+ ion = 11 p+ and
10 eCations are named the same as the metal.
sodium
Na
Na+ + 1esodium
ion
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The Periodic Law and Periodic Table
Dimitri Mandeleev ordered elements by atomic mass
Periodic Law—When the elements are arranged in
order of increasing atomic mass, certain sets of
properties recur periodically.
put elements with similar properties in the same
column
used pattern to predict properties of undiscovered
elements
Where atomic mass order did not fit other properties,
he reordered by other properties.
Te & I
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Periodic Law
13
Periodic Table
Periodic Table - Tabulates all known elements in
order of increasing atomic number.
A period consists of the elements in one
horizontal row of the periodic table.
A group consists of the elements in any one
column of the periodic table.
The groups are usually numbered.
Main group The eight “A” groups are called main
group (or representative) elements. (See Figure
2.11)
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The Periodic Table
Transition elements The “B” groups are called
transition elements.
Inner transition elements. The two rows of
elements at the bottom of the table
Elements in any one group have similar
properties.
The elements in group IA, often known as the
alkali metals, are soft metals that react easily with
water.
The group VIIA elements, known as the halogens,
are also reactive elements
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Major Divisions of Periodic table
Metals, nonmetals and metalloids
A metal is a substance or mixture that has a
characteristic luster and is generally a good
conductor of heat and electricity.
A nonmetal is an element that does not exhibit
the characteristics of the metal.
A metalloid, or semi-metal, is an element having
both metallic and nonmetallic properties.
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Patterns in Metallic Character
= Metal
= Metalloid
= Nonmetal
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The Modern Periodic Table
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Important Groups—Hydrogen
Nonmetal
colorless, diatomic gas
very low melting point and density
reacts with nonmetals to form molecular compounds
HCl is acidic gas
H2O is a liquid
reacts with metals to form hydrides
metal hydrides react with water to form H2
HX dissolves in water to form acids
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Alkali Metals
Alkali metals: Group 1A metals Na, Li, K and Rb ,
soft, low melting points, low density, all are very reactive
react violently with water to form basic (alkaline) solutions and
H2(releases a lot of heat)
2 Na + 2 H2O
2 NaOH + H2
Alkali Earth Metals: Group 2A, harder, higher melting, and
denser than alkali metals
less reactive than corresponding alkali metal
form stable, insoluble oxides from which they are normally
extracted
alkaline earth oxides are basic
reactivity with water to form H2
Be = none; Mg = steam; Ca, Sr, Ba = cold water
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Halogens
Halogens: Group 7A elements, nonmetals
F2 and Cl2 gases; Br2 liquid; I2 solid
all diatomic, very reactive
Cl2, Br2 react slowly with water
Br2 + H2O
HBr + HOBr
react with metals to form ionic compounds
HX all acids
HF weak < HCl < HBr < HI
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Noble Gases
Noble Gases- Group 8A, all gases at room
temperature
very low melting and boiling points
very unreactive, practically inert
very hard to remove electron from or give an
electron to
Argon, a small component of the earth’s
atmosphere
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Ions and the Periodic Table
The charge on an ion can often be determined from
an element’s position on the Periodic Table
Metals always form positively charged cations
For many main group metals,
the charge = the group number.
Nonmetals form negatively charged anions
For main group nonmetals,
the charge = the group number-8
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Predictable Ion charges
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What is the charge on each of the
following ions?
potassium cation
sulfide anion
calcium cation
bromide anion
aluminum cation
K+
S2
Ca2+
Br
Al3+
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Atomic Mass
We generally use the average mass of all an
element’s atoms found in a sample in calculations
However, the average must take into account the
abundance of each isotope in the sample
Atomic mass- The average mass of an elements
atoms
At. Wt. = mA x(its fractional abundance)+ mB x (its fractional abundance)
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Example 2.5 If copper is 69.17% Cu–63 with a
mass of 62.9396 amu and the rest Cu–65 with a
mass of 64.9278 amu, find copper’s atomic mass
Given: Cu–63 = 69.17%, 62.9396 amu
Cu–65 = 100 - 69.17%, 64.9278 amu
Find:
atomic mass, amu
Conceptual
isotope masses,
avg. atomic mass
Plan:
isotope fractions
Relationships:
Solution:
Check:
The average is between the two masses,
closer
the major isotope.
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18
Counting Atoms by Moles:
Mole (A chemist’s dozen): One mole of an element is
the amount of that element that contains Avogadro’s
number(6.022x 1023) of atoms.
1 mole = 6.022 × 1023 units
like 1 dozen = 12 things
One mole is equal to the number of atoms is equal to
the number of atoms in 12 g of pure C-12
1 atom of C-12 weighs exactly 12 amu.
1 mole of C-12 weighs exactly 12 g.
1 mole of C atoms =12.01 g = 6.022 × 1023 C atoms
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Example 2.6 Calculate the number of atoms in
2.45 mol of copper
Given:
Find:
Conceptual
Plan:
Relationships:
2.45 mol Cu
atoms Cu
mol Cu
atoms Cu
1 mol = 6.022 × 1023 atoms
Solution:
Check:
Since atoms are small, the large number of
atoms makes sense.
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Converting between Mass and amount
(Moles)
Molar mass-The mass of one mole of atoms of an
element
The molar mass of an element, in grams, is
numerically equal to the element’s atomic mass, in
amu.
1 mole O = 15.999 g O = 15.999 amu = 6.022x 1023
O atoms
1 mol C = 12.01 g C = 6.022x 1023 O atoms
Conversion factor
1 mol O/16.0 g O or 16.0 g O/ 1 mol O
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Sample Problem-Mass to moles
Calculate the moles of carbon in 0.0265 g of pencil
lead.
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Learning objectives
To know Modern Atomic theory
Using Law of definite proportions and multiple proportions
Working with atomic numbers, mass numbers and isotope
symbols
Determining atomic weight from isotopic masses and fractional
abundances
Predicting the charge of the ions
Calculating atomic mass, converting between mole and number
of atoms
Using mole concept and converting between mass and moles
and us
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20
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