Download Chapter 2: Atoms Molecules and Ions

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Chapter 2: Atoms and Elements
I) Imaging & Moving Atoms
1) How small can we go? (A philosophical exercise).
2) Atoms can now be seen using scanning tunneling microscopy (STM).
3) All matter is composed of atoms. The word atom comes from the Greek word
atomos meaning indivisible.
4) Atoms are the key to connecting macroscopic and microscopic worlds.
5) Atom is smallest identifiable unit of an element. There are about 91 elements
found in nature and around 20 synthetic elements.
II) Early History of Atoms & Scientific Revolutions
1)
2)
Although chemistry can be traced back to the dawn of humanity; it’s a
subject that developed globally (Egypt, China, Greece, etc.).
Greeks were the first to attempt to explain the nature of matter. They
relied on logic & reason to explain the natural world.
i) Aristotle & Plato believed that there were four elements (air,
water, Earth, & fire).
ii) Leucippus & Democritus proposed that elements are composed
of tiny particles that we call atoms.
iii) No way to establish which idea was correct; public taken in by
celebrity of Aristotle & Plato.
3)
4)
Scientific revolution in 16th century brought rapid advancement in science
by developing scientific method.
Robert Boyle was the first to study chemistry through rigorous chemical
experiments. Authored the book The Sceptical Chemist where he argued
against the Greek notion of studying nature. Boyle was also the first to
define an element.
III) Fundamental Chemical Laws: The Road to Atomic Theory
1) Law of Conservation of Matter (Mass) (A. Lavoisier)
A) Matter can neither be created nor destroyed, but can be converted from
one form into another.
B) Lavoisier typically called The Father of Modern Chemistry.
2) Law of Definite Proportions
(J. Proust)
A) A given compound always contains EXACTLY the same proportion of
elements by mass.
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3) Law of Multiple Proportions
(J. Dalton)
A) Elements can combine in different ways to form different substances,
whose mass ratios are small whole number multiples of each other.
Example
Cmpd. I
Cmpd. II
1.33 g O / 1.00 g C
2.67 g O / 1.00 g C
CO
CO2
4) Dalton’s Atomic Theory (1808)
A) Four Postulates (go over these)
B) Dalton prepared the first table of atomic weights (atomic masses)
Many entries were wrong due to assumption that atoms combine in
simplest fashion.
Example:
water
water
OH (Dalton)
HOH (modern)
8:1
16:2
(8:1)
IV) Early Experiments to Characterize the Structure of the Atom
1) J.J. Thompson
A) Worked with cathode ray tubes.
B) Discovered the electron and surmised that all atoms contain electrons.
C) Proposed the first model of the atom (named “plum pudding model of
atom” after his favorite dessert.
D) Calculated the charge to mass ratio of the electron.
e/m = -1.76 x 108 C / gram
2) Robert Millikan
A) Determined charge of an electron (e- = 1.602 x 10-19 C)
B) Allowed one to calculate mass of electron (me =. 9.11 x 10-31 kg)
3) Henri Becquerel
A) Discovered radioactivity
B) Three types of radioactive particles
i)
ii)
iii)
alpha particles have a 2+ charge and a helium nucleus
beta particles are merely electrons
gamma rays are a highly energetic form of radiation
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4) Ernest Rutherford
A) Tested and modified “plum pudding model of atom”
B) Emphasize scattering of alpha particles (CAUSED ONLY BY
CENTER OF CONCENTRATED + CHARGE)
V) Modern View of Atomic Structure
1) Atomic Structure
A) nucleus
(protons and neutrons)
i) majority of mass of atom found in tiny, highly dense place.
B) electron cloud
(electrons)
i) volume of atom found here.
ii) chemistry of atoms occurs because of interactions between
electron clouds.
C) Table of atomic constituents
Atomic Constituent
electron (e-)
proton (p+)
neutron (no)
Mass of Component
9.11 x 10-31 kg
1.67 x 10-27 kg
1.67 x 10-27 kg
Charge
11+
0
2) Atomic Mass
A) Atomic mass unit (amu) is defined as exactly 1/12th the mass of a
carbon 12 atom containing six protons and six neutrons.
B) Expressed in this unit, the mass of proton and neutron is approximately
1 amu. The mass of the electron is negligible and can be ignored.
3) Atomic Number & Identifying Elements
A) The number of protons in an atom’s nucleus defines the element.
B) Atomic Number symbolized by the letter Z.
C) Each element in the periodic table is symbolized by a chemical symbol.
The chemical symbol for an element is written with the first letter always
capitalized and subsequent letters in lowercase.
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Examples:
S
Mg
Cu
Co
Sulfur
Magnesium
Copper
Cobalt
(From Latin cuprum)
While many of the chemical symbols are easily related to the name of the
element; some are derived from Latin or other languages. These
exceptions need to be learned (see Figure 2.10 in text).
Each student is responsible for knowing the name and chemical
symbol for the first 36 elements for Exam I.
Chemical symbols are very important since they are the backbone of
chemistry. For example, Co is the element cobalt while CO is carbon
monoxide, a gaseous compound that is found in air pollution!
4) Isotopes
A) Atoms with same number of protons (#p+) but different numbers of
neutrons (#no). Isotopes have the same atomic number (Z) but different
mass numbers.
B) Recognizing Isotope Notations
i) X = element from periodic table
ii) A = mass number = #p + #n
iii) Z = atomic number = #p
-NOTE: In a neutral atom, Z = #p = #e
C) Examples
Na-23
Sodium-23
(Z=11)
11 p, 12 n, 11 e
Na-24
Sodium-24
(Z= 11)
11 p, 13 n, 11 e
5) Ions
A) Ions are charged atoms.
B) Charges arise due to loss or gain of electrons.
C) If atom loses electrons, a positively charged cation is formed.
Example: Li ──> Li+ + 1 e-
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D) If atom gains electrons, a negatively charged anion is formed.
Example: F + 1 e- ──> FE) In ordinary matter, cations and anions always occur together so
that matter is charge-neutral overall.
VI) Elements & The Periodic Table
1) The development of the periodic table can be traced back to the work of
scientists like Lavoisier, Dobreiner, Newlands, Meyer & Mendeleev.
2) When elements are arranged in table, similar physical & chemical
characteristics repeat at regular intervals, this phenomenon is called
periodicity.
3) The elements are arranged in the periodic table in seven rows, also termed
periods. Within each column, called a group or family, similar physical and
chemical characteristics are found. There are 18 columns in the periodic table.
4) Be familiar with the following Sections of the Periodic Table.
A) Main Group Elements
B) Transition Metals
C) Inner Transition Metals
(Groups IA to VIIIA)
(Center of Table)
(Bottom Two Rows: known as the
Lanthanides and Actinides).
5) Types of Elements in Table
A) Metals
- most common type of element
- found on left side of table
- conduct electricity, malleable, ductile, high mp, high bp
- most are solids at room temperature
- metals tend to lose electrons when undergoing chemical changes
B) Nonmetals
- found in upper right portion of table
- nonconductors, lower mp and bp vs. metals
- most are liquids or gases at room temperature
- nonmetals tend to gain electrons during chemical changes
- form covalent compounds with one another by sharing electrons
C) Metalloids (semi-metals)
- share properties of both metals and nonmetals depending on
chemistry involved.
- useful industrially (Si, As, Ge, etc.).
- found along dark diagonal line in table
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6) Identifying Elements within Periodic Table
A) Rows
periods
B) Columns groups or families
Group IA
Alkali Metals
Group IIA
Alkaline Earth Metals
Group IIIA - VIA
Name of 1st member + Family
Group VIIA
Halogens
Group VIIIA
Noble Gases (Inert Gases)
C) Be familiar with general properties of main group elements (see text).
7) Ions & Periodic Table
A) A main-group metal tends to lose electrons, forming a cation with the
same number of electrons as the nearest noble gas.
B) A main-group nonmetal tends to gain electrons, forming an anion with
the same number of electrons as the nearest noble gas.
C) Review Figure 2.14.
VIII) Atomic Mass
1) Atomic Mass of Element represents the weighted average of all the naturally
occuring isotopes of that element.
2) Value found below the atomic symbol of the element in the periodic table.
A) Calculating Atomic Mass for an Element
i)
ii)
iii)
iv)
Must be given relative abundance (%) for each isotope
Must know isotopic mass for each isotope
Convert percent abundance to a decimal (divide by 100%)
Use following formula to obtain result
Atomic Mass = Σ (fraction of isotope n) x (mass isotope n)
Example:
Find Atomic Mass of Carbon given the following data.
Data:
98.89% Carbon 12
1.11% Carbon 13
isotopic mass 12 amu
isotopic mass 13.0034 amu
Solution: (12 amu * 0.9889) + (13.0034 amu * 0.0111)
= 11.867 amu + 0.144 amu = 12.011 amu
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IX) Avogadro’s Number & The Mole
1) Since molecules are so small, it is necessary to convert a number ratio of
reactant molecules into a mass ratio for a chemical reaction to be useful.
2) Mass ratios are determined by using atomic masses for the elements.
i) Atomic masses (atomic weights) are found in periodic table beneath the chemical
symbol, and represent the average of all the naturally occurring isotopes of that
element.
3) For more complicated substances (molecules, ionic compounds, etc.) one
makes use of the formula mass (molecular mass).
i) Formula mass (FM) represents the sum of atomic masses of all atoms
in a formula unit of any compound (ionic or molecular).
ii) Molecular Mass (MM) is the sum of the atomic masses for all atoms
in a molecule.
iii) In some texts, one will see the terms formula weight (FW) and
molecular weight (MW). These have the same definitions as above.
iv) Examples
HCl
1.01 amu + 35.45 amu = 36.55 amu
NaCl 23.00 amu + 35.45 amu = 58.45 amu
1) How are molecular (formula) masses used?
i) Samples of different substances contain the same number of molecules
(or formula units) whenever their mass ratio is the same as their
molecular-mass (or formula-mass) ratio.
ii) A particularly convenient way to use this mass-number relationship is
to measure amounts in grams that are numerically equal to molecular
masses.
iii) Chemists have adopted a special unit called a mole. The mole is the SI
unit for the amount of a substance.
iv) A mole is defined as the amount of a substance that contains as many
“entities” as there are in exactly 12 g of carbon-12. Other more useful
relationships based on the mole include the following.
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1 mole of X = NA units
(units = atoms, ions, molecules)
NA = Avogadro’s Number = 6.022 x 1023 units/mole
mass of 1 mole of a substance X = molar mass of X in grams.
Molar mass is equal to formula (molecular) mass of X in
grams. Molar mass values are expressed in grams/mole.
v) Example Exercises
Given:
Unknowns:
Rel. Info:
10.0 g Al
? moles Al
? Al atoms
26.98 g Al = 1 mole Al
(molar mass of Al)
Solutions:
10.0 g Al x (1 mole Al / 26.98 g Al) = 0.371 mole Al
0.371 mole Al x (6.022 x 1023 atoms Al / 1 mole Al) = 2.23 x 1023 atoms Al
Check: Answers have corresponding units and are reasonable values.
Given:
Unknowns:
5.00 x 1020 atoms Co
? moles Co
? g Co
Rel. Info:
58.93 g Co = 1 mole Co
1 mole Co = 6.022 x 1023 atoms Co
Solutions:
5.00 x 1020 atoms Co x ( 1 mole Co / 6.022 x 1023 atoms Co) = 8.30 x 10-4 moles Co
8.30 x 10-4 moles Co x ( 58.93 g Co / 1 mole Co) = 4.89 x 10-2 g Co = 0.0489 g Co
Check: Answers have correct units and seem plausible.
Given:
Unknown:
Rel. Info:
1.56 x 10-2 g sample Juglone C10H6O3
? MM of Juglone
? moles of Juglone in sample
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Note: Here is the solution to the first part. I am placing it in under relevant
information because we will use it to solve part two of the problem.
C10H6O3
10 C
10 x 12.01 g/mole
6H
6 x 1.008 g /mole
3O
3 x 16.00 g/mole
MM of Juglone
120.1 g/mole
6.048 g/mole
48.00 g/mole
174.1 g/mole
Solution:
1.56 x 10-2 g C10H6O3 x ( 1 mol C10H6O3 / 174.1 g C10H6O3 )= 8.96 x 10-5 moles
C10H6O3
Check: Answer has correct units and seems reasonable.
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