Download Atoms, Molecules and Ions

Atoms, Molecules and Ions
Chapter 2
In This Chapter
• History of atoms.
• Subatomic particles.
• Atomic numbers, mass numbers.
• Isotopes and Atomic weights.
• Compounds, Molecules and Ions.
• Nomenclature.
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Elements
• An element is a pure substance that can’t be
changed into a simpler form of matter by a
chemical reaction.
• Elements consist of atoms; the smallest structural
unit of an element that retains the chemical
properties of that element.
• Atoms are considered as the building blocks of
every chemical around us.
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The Origin of Elements
• After a very short time (matter of seconds) that the universe was
produced by the Big Bang the only elements present were hydrogen
and helium.
• After millions of years the cooling of the universe caused the
hydrogen and helium to collect together and form large clouds, due
to gravity.
• These clouds eventually contracted, became hotter and hotter and
burst to become stars.
• Intense heat within the stars made the hydrogen and helium atoms to
smash together and produce other heavier elements.
• This merging releases even more heat which is responsible for
starlight.
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Names of Elements
• Names come from various sources including early
Greek, Latin or German names describing some
property of the element, names of places, names of
people.
– Iodine from Greek word iodes – violet like.
– Bismuth from German weisse masse – white mass,
miners called it wismat.
– Germanium and Francium – discovered in
– Element number 110 was recently named
Darmstadtium to honor the German city in which it was
discovered.
– Einsteinium and Curium – to honor
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Symbols of Elements
• An universally accepted abbreviation for the name
of the element.
• Elemental symbols are unique.
• 14 elements have single letter symbols, while the
rest have 2.
• The following rules are typically followed:
– Symbols have either one or two letters.
– If only one letter is used it is capitalized, C, H, O.
– If 2 letters are used only the first letter is capitalized,
Na, Ca, Hg.
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Early Atomic Theories
• Democritus, a Greek philosopher (460 –370 BC)
proposed the idea of an “atom” as an indivisible
component of matter, using the word atomos
which means indivisible.
• John Dalton, in 1808, published A New System of
Chemical Philosophy in which he presented his
theory of atoms.
• Dalton’s theory was the first one to provide a
systematic description of the relationship between
atoms and matter.
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Dalton’s Atomic Theory
• Dalton’s basic postulates (ideas) were:
– All matter is composed of atoms which are indivisible
and indestructible. Atoms are considered as the
ultimate chemical particles.
– An element is composed of identical atoms with fixed,
identical properties and masses.
– Compounds are formed by the combination of atoms
of 2 or more different elements in a fixed whole
number ratio.
– A chemical reaction involves a combination,
separation or rearrangement of atoms. Atoms are
neither created nor destroyed.
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Problems with Dalton’s Theory
• Atoms can be divided further.
• Elements can have more than one mass.
• Here is what we know about the atom now:
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Regions of an Atom
• All atoms consist of 2 distinct regions.
• The nucleus: A small dense region at the center
which contains positively charged protons and
neutral (with no charge) neutrons.
• Surrounding the nucleus is a diffuse region which
contains the negatively charged electrons.
• Atoms of different elements have different # of
protons, electrons and neutrons.
• The # of protons decides the identity of the
element.
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Properties of Subatomic Particles
Proton (p)
Charge = 1 Mass = 1.6725 x 10-24 g
Neutron (n) Charge = 0 Mass = 1.6750 x 10-24 g
Electron(e-)
Charge =
-1
Mass = 9.1095 x 10-28 g
• Protons and neutrons have masses that are similar to
each other.
• The electron is about 1800 times lighter than the other
2.
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The Neutral Atom
• Neutral = No charge
• # of protons (positive charges) = # of electrons
(negative charges)
• Atoms of different elements have different
numbers of protons, neutrons and electrons.
• The number of protons determine the identity
of the element.
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4
Atomic Symbol
Atomic Mass Number
(A) = The sum of # of
protons and the # of
neutrons
Atomic number (Z) =
The # of protons
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C
In any neutral atom, the atomic number = # of electrons in
one atom of that element
Atomic Mass Number – Atomic number = Number of
neutrons.
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Isotopes
• Dalton, had proposed that all atoms of the same
element must have the same mass.
• One of the 2 flaws in Dalton’s theory was, atoms of
the same element can have DIFFERENT masses.
• The mass of an atom is due to the mass of the
protons and the neutrons in that atom.
• Isotopes are atoms of the same element which have
the same number of protons but different number
of neutrons.
The word isotopes is from the Greek words for “equal place”.
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Isotopes
• Isotopes have the exact same physical properties but
different chemical properties. (Some are
RADIOACTIVE and can emit particles and energy).
• All naturally occurring elements have 2 or more
isotopes.
• Isotopes exist in different amounts (called isotopic
abundance) and have different lifetimes.
• A mass spectrometer is the most direct and accurate
means of determining atomic weights and the
existence of isotopes.
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Example of Isotopes
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C
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6
C
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6
C
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Atomic Mass
• The mass of an atom is decided by the number of
protons and the number of neutrons present in the
atomic nucleus.
• Since atoms are so tiny, using an ordinary unit
such as gram or kilogram is inconvenient and the
atomic masses are usually reported in a special
unit called atomic mass unit (amu).
• One amu = 1.661 x 10-24 g
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How and Why of the amu
• Dalton had proposed an atomic scale based on the
mass of one Hydrogen atom, which he had
assigned the mass of 1. Several of his reported
masses were later proven wrong.
• The Hydrogen based scale was then replaced with
an Oxygen based scale and then by the now
accepted Carbon – 12 based scale.
• One amu = 1/12th of the mass of one carbon atom.
• The mass of one proton = 1 amu.
• The mass of one neutron = 1 amu.
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Atomic Masses and Isotopes
• The sum of the atomic masses of the protons and
the neutrons present in an atom should add up to
give an exact round number and be the same as
the atomic mass number “A”.
• The presence of isotopes makes matters slightly
complicated.
• Each isotope of an element is present in certain
percent abundance, which has to be accounted for
when it comes to calculating the atomic masses.
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Using % abundance
• The % abundance is factored in the calculation
for finding the atomic mass of any element.
• Sample Exercise 2.4 shows how to make use of
this idea.
• Consider the following example:
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Isotopes of Neon
• Neon has 3 naturally occurring 1. Change each % value to a
isotopes. The masses and
fraction.
natural abundances of these
isotopes are as follows.
Calculate the atomic mass of
Neon.
1 19.99 amu 90.48 %
2 20.99 amu 0.27 %
3 21.99 amu 9.25 %
2. Multiply the isotopic masses by
these fractions and then add the
results together.
Atomic mass of Neon = (Mass of
isotope #1 x fraction of #1) +
(Mass of isotope #2 x fraction of
#2) + (Mass of isotope #3 x
fraction of #3)
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Thus, the average atomic mass of Neon is 20.18
amu
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Compounds, Molecules & Ions
Compounds
Molecular
Ionic
Molecules composed of 2
or more elements
Cations with
positive charges
Anions with
negative charges
There are more than 19 million registered compounds with no end
in sight as to how many more will be prepared.
Each compound can be identified and distinguished by its
characteristic properties.
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The Periodic Table
• A systematic arrangement of elements according to increasing atomic numbers.
Comic book periodic table (http://www.uky.edu/Projects/Chemcomics/)
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The Periodic Table
• A systematic arrangement of elements according to increasing atomic numbers.
• The vertical columns are known as groups.
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The Periodic Table
• The horizontal rows are known as periods.
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Metals, Metalloids and Non-metals
• Metals:
– Solids at room temperature
(except Hg).
– Hard, shiny.
– Good conductors of
electricity and heat.
– Malleable (from Latin for a
hammer).
– Ductile.
– High melting point and
density.
– Typically located on the left
side of the periodic table.
– Metals will typically
combine with non-metals to
form ionic compounds.
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Metalloids
• Have properties
that are in between
the metals and
non-metals.
• Certain metalloids
are used as the raw
materials for the
manufacture of
semiconductor
devices.
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Non-metals
• Solids, liquids (Br) or
gases.
• Dull, non-malleable
and non-ductile.
• Low melting points
and densities.
• Poor conductors of
heat and electricity.
• Located on the right
of the periodic table.
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Elements that exist as diatomic molecules
• There are 7 known elements that exist as diatomic
molecules:
H2, F2, O2, N2, Cl2, Br2, I2
• When any of these elements are mentioned it is
assumed that they exist as diatomic molecules
unless specified otherwise.
• 2 other elements that are commonly polyatomic are Phosphorus (P4) and Sulfur (S8)
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Chemical Formulas
• Show the symbols and the ratios in which the atoms of the
elements in a compound combine.
• Do not show the arrangement of atoms in the compound
(structural formulas).
• If the formula contains one atom of an element the symbol
of the element represents the atom. The number 1 is not
used as a subscript. The formula of water is H2O, not
H2O1.
• When the formula contains more than one atom of an
element the number of atoms is indicated by a subscript
written to the right of the atom. The number 2 to the right
of the H indicates 2 atoms of hydrogen present in H2O.
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Chemical Formulas (continued)
• When the formula contains more than one of a group of
atoms that occurs as a unit, parentheses are place around the
(group) and the number of units of the group are indicated
by a subscript written to the right of the parentheses.
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Ions
• Ions are formed by the loss (cations) or gain
(anions) of electrons by neutral atoms.
• Cations: Positively charged ions that are formed
an element loses one or more electrons.
(Table 2.4 in the book provides some examples)
Atom
Cation
K Potassium
K+ Potassium ion
Ca Calcium
Ca2+ Calcium ion
Al Aluminum Al3+ Aluminum ion
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Ions (continued)
• Anions: Negatively charged ions that are formed
when a neutral atom gains one or more electrons.
• Anions are named by using the stem of the parent
element and changing the ending to -ide.
(Table 2.5 in the book provides some examples)
Atom
Cl Chlorine
Br Bromine
O Oxygen
N Nitrogen
Anion
Cl- Chloride ion
Br- Bromide ion
O2- Oxide ion
N3- Nitride ion
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Ionic charges and the periodic table
• With groups 1A and 2A the charge on the cations
= group number.
• With groups 5A, 6A and 7A the charge on the
anions = 8 – group number
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Name to Formula: Ionic Compounds
• Ionic compounds = Simplest possible combination
of cations and anions.
• All compounds have a net electric charge = zero.
• The charges on the cations and the anions should
cancel each other out.
• Consider sodium chloride:
– Made of the sodium ion (Na+) and the chloride ion (Cl-)
– The formula for this compound would be NaCl.
– The following formulas will be wrong:
Na2Cl2, NaCl2, Na2Cl.
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More Ionic Compounds
• Write the formula for Calcium fluoride: Calcium
ion is Ca2+, fluoride is F-, only possible
combination is CaF2.
• Aluminum oxide: Aluminum:Al3+, oxide O2-:
Al2O3 (The least common multiple of 2 and 3 is 6,
where 2(3+) + 3(2-) = 0)
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Even More Ionic Compounds
Name
Ions
Lowest
common
multiple
Sum of charges
on the ions
Potassium
iodide
Potassium
sulfide
Copper
sulfate
K + & I-
(+1) + (-1) = 0
K+ & S2-
2(+1) + (-2) = 0
Cu2+ &
SO42-
(+2) + (-2) = 0
Ammonium
phosphate
Aluminum
chromate
NH4+ &
PO43Al3+ &
CrO42-
Formula
3(+1) + (-3) = 0
2(3+) + 3(2-) =
0
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Metal cation - monatomic nonmetal anion
• General formula MaXb.
• Always write the name of the metal ion followed
by the name of the anion.
• Need to indicate the charge on the cation except
when metals from groups 1A and 2A are present or
the ammonium ion is present.
• NaCl
• MgO
• Fe2O3
• CuCl2
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Metal cation - polyatomic nonmetal anion
• General formula: MaXb
• Same rules as before.
• Only difference is the anion contains more than
one elements.
• Na2SO4
• Ca3(PO4)2
• K2Cr2O7
• (NH4)3PO4
• ZnCO3
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Molecular compounds
• Nonmetal-nonmetal compounds.
• Look for the absence of common metals.
• General formula: AxBy
– Write the numerical prefix for nonmetal A. If there is only one
atom of A present there is no need to use “mono”. If A is
hydrogen then a numerical prefix is not used.
– Write the full name of A.
– Write the numerical prefix for y, if y = 1, no need to write mono.
– Add the stem of the name of the nonmetal B.
– Complete the name by adding the suffix –ide to the stem of B.
• CO2, N2O5, CCl4, SO3.
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Acids
• An important class of compounds that can be considered
as molecules in which one or more H+ ions are attached to
an anion.
• If the anion does not contain oxygen the acid is named by
the prefix hydro- and the suffix –ic.
– HCl: Hydrochloric acid & HF: Hydrofluoric acid
• If the anion ends with an -ate the suffix –ic is added to the
root name.
– H2SO4: Sulfuric acid (SO42- sulfate ion), H3PO4: Phosphoric
acid.
• If the anion ends with an –ite, the suffix –ous is added to
the root name.
– H2SO3: Sulfurous acid (SO32- sulfite ion), HNO2: Nitrous acid
(NO2- nitrite ion).
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Organic compounds
• Contain primarily C
and H but may also
have N, O, S and P.
• Alkanes: Simplest kind,
contain only C and H in
the ratio CnH2n+2.
• Methane (CH4), ethane
(C2H6), Propane (C3H8)
and butane (C4H10).
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