Download Chapter 2 – Elements

Chapter 2 – Elements
 In ancient times, symbols for heavenly bodies were used for the 7
elements that were known at that time.
 In the 1800’s More elements were discovered and John Dalton
(English)devised a new series of symbols for all the elements.
 In 1814 J.J. Berzelius (Swedish) devised a system of using letters as
symbols for the elements.
 His system of symbols has 4 characteristics:
1) international 2) precise 3) logical
4) simple
 There is a governing body (IUPAC – International Union of Pure
and Applied Chemistry) which sets down the rules for naming
compounds and elements. They still use Berzelius’s symbols.
IUPAC rules
1. Symbols are international.
2. The first letter of the symbol is upper case.
3.The name of any new element ends in “ium”.
4. Elements beyond 103 have temporary names and symbols
derived by combining (in order)the “element roots” for the atomic
number and the adding “ium”.
Number
0
1
2
3
4
Root
nil
un
bi
tri
quad
Metals and Nonmetals
Number
5
6
7
8
9
Root
pent
hex
sept
oct
enn
 When stating the properties of elements, it is important for chemists
to specify the conditions for which the elements have those specific
properties.
 Unless other wise stated, the standard set of conditions is called
SATP – standard ambient temperature and pressure.
 SATP – 25 oC and 100. kPA
 There is another set of standards, STP, that is sometimes used.
 STP – 0.00oC and 101.325 kPA
Metals
shiny/luster
flexible (malleable,ductile)
solids (except Hg)
good conductors of heat and
electricity
left of the staircase
form +ve ions
Nonmetals
dull
brittle
most are gases (over half)
poor conductors
right of the staircase
form –ve ions
Metalloids
Solids, high melting points, non-conductors or semi-conductors
fall on the staircase, form +ve and –ve ions
The Periodic Table
As more and more elements were discovered in the early 1800’s
scientists searched for a way to organize and classify the elements.
1. J. Dobereiner (German – 1817)
 discovered that several groups of 3 elements had similar properties.
 called these groups triads. ie. Ca Sr Ba
Cl Br I
2. John Newlands (English – 1863)
 arranged elements by increasing atomic mass.
 came up with the periodic law – When elements are arranged in
order of increasing atomic mass, chemical and physical properties
form patterns that repeat at regular intervals.
 initially he called this law “The Law of Octaves” because properties
repeated every eighth element.
3. Dimitri Mendeleyev (Russian- 1869)
 ideas similar to Newlands but properties repeated after periods of
varying length.
 arranged elements by increasing atomic number and not mass.
 had blanks in his table and he predicted that elements would be
discovered that would fit into these empty spaces. (even predicted
their properties)
alkali metals
group 1, soft metal, silver colored, react violently with water, all react
with the halogens
alkaline earth metals
group 2, light in mass, very reactive metals, readily react with oxygen to
form oxides
halogens
group 17, extremely reactive, F is the most reactive, all react with
hydrogen
noble gases
group 18, very stable and non-reactive
transition metals
groups 3 – 12, exhibit a wide range of physical and chemical properties
representative elements
groups 1,2 and 13 – 18
row/period/series (horizontal) – elements properties gradually change
from metallic to nonmetallic.
group/family/column (vertical) – elements having similar chemical
properties.
Atomic Theories
1. Greek Philosophers (500 B.C.)
 believed all substances to be composed of small individual particles
called atoms.
 severely criticized by Aristotle.
2. Aristotle’s Theory (384 – 332 B.C.)
 all matter is made up of 4 basic substances earth, air, fire, water.
 eventually many predictions and explanations using Aristotle’s
theory were shown to be false.
 this led to the revival of the atom concept.
3. J. Dalton (1803) Billiard Ball Theory
 all matter is composed of tiny, indivisible particles called atoms.
 atoms of an element have identical properties.
 atoms of different elements have different properties.
 atoms combine in constant ratios to form new substances.
Dalton was successful in explaining the following laws :
a) Law of Conservation of Mass
b) Law of Definite Composition
c) Law of Multiple Proportions
4) J.J. Thompson (1897) Raisin Bun/Blueberry Muffin
Theory
 found negatively charged particles called electrons using a cathode
ray tube.
 his model of the atom had electrons distributed evenly inside a
spherical positively charged atom.
5) H. Nagaoka (1904) Saturn Theory
 he suggested that the atom was a positively charged sphere
surrounded by a ring of electrons.
 No evidence until 1911 to contradict Thompson’s or Nagaoka’s
theory.
R. Milikan
 found the charge on an atom to be -1 and have a mass 1/1837 of the
total mass of the hydrogen atom.
6) E. Rutherford (1914) Nuclear / Beehive Model
 through a gold foil experiment he found out that most of the atom
was empty space with a small positive nucleus.
 called the positive particles protons.
 proposed that the electrons moved around the nucleus like bees
around a beehive.
H.G.J. Moseley
 first to show the relationship between the atomic number and
nuclear charge.
J. Chadwick
 discovered neutral particles in the nucleus which he called
neutrons.
 neutrons have the same mass as protons.
F. Soddy
 first to propose the idea of isotopes
 isotope – atoms having the same number of protons but varying
numbers of neutrons.
7) N. Bohr (1921) Planetary Model
 said electrons moved around the nucleus in circular fixed orbits.
 each electron had a fixed amount of energy related to its orbit.
 electrons can move to unfilled orbits if a quantum of energy is
absorbed or released.
 the further an electron is from the nucleus, the more energy it has.
 an atom with a filled outer energy level is stable. (unreactive)
8) Quantum Mechanical Theory
 uses 4 quantum numbers to describe the behavior of electrons.
i)Principle Quantum Number (n)
 this refers to the level or shell of the electron. (general region)
 the maximum # of electrons in a shell = 2n2 (n = shell number)
1st shell (k)
2nd shell (l)
3rd shell (m)
4th shell (n)
2(1)2
2(2)2
2(3)2
2(4)2
=
=
=
=
2 electrons
8 electrons
18 electrons
32 electrons
ii) Secondary Quantum Number ( )
 also called the azimuthal quantum number
 refers to the sublevel or subshell of the electron
Shell
1
2
3
4
# of Subshells
1
2
3
4
Symbols
s
s,p
s,p,d
s,p,d,f
iii) Magnetic Quantum Number (m)
 refers to the orbital, direction or shape of the electron cloud
 each orbital can hold 2 electrons.
Subshell
s
p
d
f
# orbitals
1
3
5
7
# of electrons
2
6
10
14
iv) Spin Quantum Number
 refers to the spin on the electrons
 each electron in an orbital has either a clockwise spin(+ ½ ) or a
counterclockwise spin (- ½)
Pauli Exclusion Principle – no two electrons in an atom can have the
same 4 identical quantum numbers.
Hund’s Rule – electron’s are put into orbitals one at a time before they
are paired up.
Electrons fill up the 1st level before going on to the 2nd level etc.
Electrons fill up the s sublevel before going on to the sublevel etc.
Electron Configurations
This a description of the energy level and sublevel for all electrons in an
atom.
ie.
H
Be
O
= 1s1
= 1s2 2s2
= 1s2 2s2 2p4
Noble Gas Notation
This is a shortcut for doing the electron configuration.
1. Write the symbol for the noble gas in the row which precedes the
row in which the element is in.
2. Put square brackets around the symbol for the noble gas.
3. Continue the electron configuration from the noble gas to the
element.
ie. Tc = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d5
[Kr] 5s2 4d5
Nd = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14
[Xe] 6s2 4f14
Orbital Notation
Write out the electron configuration or the noble gas notation and then
use boxes or circles to represent the orbitals and arrows to represent the
electrons.
* Place 1 electron in each orbital of a sublevel before pairing them
up.
ie. N = 1s2 2s2
2p3
Cu = 1s2 2s2
2p6
3s2
3p6
4s2
3d9
Diagonal Rule (Aufbau Diagram)
1s2
2s2
3s2
4s2
5s2
6s2
7s2
2p6
3p6
4p6
5p6
6p6
7p6
3d10
4d10 4f14
5d10 5f14
6d10
Lewis Diagrams
Lewis diagrams are a method to represent or show the valence electrons
(outer shell) in an atom.
Steps:
1. Write the symbol for the element. (This represents the nucleus and all
inner shell electrons.)
2. Determine the number of valence electrons.(only those
in the s and p subshell of the outermost shell)
Group Number
1
2
13
14
15
16
17
18
# of Valence Electrons
1
2
3
4
5
6
7
8
* Many of the transition elements and those in the Lanthanide
and Actinide series have 2 valence electrons.
3. Each side of the element represents an orbital. (4 sides – 1s, 3p)
3:00 represents the “s” orbital. Fill it first. Then move
counterclockwise putting one electron into each orbital before
pairing them up.
ie.
Ne
Na
Ge
Particles of the Atom
Particle
electron
proton
neutron
Symbol
e
p
n
Charge
-1
+1
0
Mass
9.11 x 10-28
1.67 x 10-24
1.67 x 10-24
Location
around nucleus
nucleus
nucleus
Atomic mass unit – the mass of a proton or a neutron.
1 u = 1.67 x 10-24 g
Atomic number (z) – equal to the number of protons in the atom.
Since an atom is neutral, the atomic number is also equal to the number
of electrons.
Mass Number (A) – is equal to the sum of the number of protons and
neutrons.
number of neutrons = A - Z
Isotopes and Average Mass
Isotope – atoms having the same atomic number but different atomic
masses. (neutrons vary)
ie. Carbon has 3 isotopes
C 12 - 6 p, 6 e, 6 n
C 13 - 6p, 6 e, 7 n
C14 - 6 p, 6 e, 8 n
The atomic mass on the periodic table is the average weighted mass.
It is calculated or based on the % of each isotope in a naturally occurring
sample.
Formation of Monatomic Ions
ion – that which is produced when atoms gain or lose electrons and gain
an electrical charge.
monatomic ions – single atoms that have gained or lost electrons.
* atoms of the representative elements form monatomic ions when
gaining or losing electrons to form the same stable electron structure of
the nearest noble gas. (filled valence shell – 8 valence electrons)
This theory of monatomic ion formation is basically restricted to the the
representative elements.
cations – positively (+ve) charged ions.
- produced by the loss of electrons.
- metallic elements form cations.
anions
– negatively (-ve) charged ions.
- produced by the gain of electrons.
- nonmetallic elements for anions.
Periodic Properties
Many properties of the elements are periodic in nature. That means
that it has a property that repeats itself at regular intervals.
There are 4 periodic properties that we will study. 1)Atomic radii 2)
Ionic radii 3) Ionization energy and 4) Electronegativity
An elements position on the periodic table and its properties are a
result of its electron configuration or number of valence electrons.
Changes in the atomic radii, ionization energy and electronegativity
can be explained based on 3 factors:
1. number of shells
2. shielding effect
3. nuclear charge
Atomic Radii
Atomic radii – is a measure of the radius of an atom.
First and foremost, as the principal quantum number increases, the
size of the electron cloud increases. Therefore as you go down a
column, the radii of the atoms increase.
As you go across a row, the size of the cloud decreases because the
nuclear charge is increasing. Even though there are more electrons,
they are being put into the same shell or a lower shell.
Ionic Radii
Ionic radii – is a measure of the radius of an ion.
Negative or nonmetallic ions have radii that are larger than their
corresponding atoms because there are more electrons being
attracted by the same nuclear charge. As a result of this, each
electron has a weaker pull on it from the nucleus and therefore the
electrons can move further away from the nucleus.
Positive or metallic ions have radii that are smaller than their
respective atoms for two reasons. When a metal loses electrons, it is
losing its valence shell, and there are also fewer electrons being
attracted by the same number of protons. This produces a stronger
pull on each electron and they are pulled closer to the nucleus.
Ionization Energy
Ionization energy – a measure of the energy needed to remove the
most loosely held electron (valence) in an atom.
In general, as you go down a column, the ionization energy
decreases.
Reason – valence electrons are further from the nucleus, plus all the
inner shell electrons shield the valence electrons from the pull of the
nucleus.
In general, as you go across a row (left to right), the ionization
energy increases because of the increase in the nuclear charge.
There are fluctuations as you go across a row. If a sublevel is full or
half-filled, the atom is more stable and therefore more energy is
needed to remove an electron from an atom.
Summary
Nuclear charge – larger the nuclear charge, the greater the ionization
energy.
Shielding effect – the greater the shielding effect, the less the ionization
energy.
Radius – the larger the radius, the less the ionization energy.
Sublevel – an electron from a full or half-full sublevel requires more
energy to be removed.
Electronegativity
Electronegativity – this is a measure of an atoms attraction for an
additional electron when the atom is in a compound.
The same factors that affect ionization energies also affect
electronegativity. Metals have low electron affinities and non-metals
have high affinities.