Unit 2
 There
are about 118 different known
elements with 88 of them occurring
naturally.
 The
names and symbols of each element
are listed on the periodic table in your
book.
 Names
for elements come from many
different sources.
• Many have Latin or Greek origins
• However, more recent discoveries are named for
descriptions, Famous Scientists, or Place of
discovery
 We
use abbreviations to simplify when
writing called Element Symbols.
 Examples:
• Berkelium (Bk)- named for place of discovery-
Berkeley, California
• Copper (Cu)- Latin- cuprum or cyprium,
discovered in Cyprus
• Lead (Pb)- Latin,- plumbum, meaning heavy
• Oxygen (O)- French, oxygene, generator of acid,
derived from the Greek, oxy and gene meaning
acid forming
 Oxygen was thought to be part of all acids, but it’s not.
Can you guess which element is part of all acids?
 Element
Symbols- First or first two letters
of the element names.
• The first letter is always capitalized
• The second letter is always lowercase
 O-
oxygen, C- carbon, Ne- neon,
Si- silicon
 Element
Symbols- Sometimes the first
two letters are not the first two letters of
the name.
• Symbols are from the old names
• Symbols are from other letters in the name
 Gold-
Aurum- Au
 Lead- Plumbum- Pb
 Zinc- Zn
 Cadmium- Cd
 Law
of Constant Composition: A given
compound always contains elements in
exactly the same proportion by mass.
 This
observation along with others
became the basis for Dalton’s Atomic
Theory.
 1. Elements
are made of tiny particles
called Atoms. All matter is made of
atoms. Atoms are indivisible and
indestructible particles.
 2. All
atoms of a given element are
identical, both in mass and in properties.
 3. The
atoms of a given element are
different from those of any other element.
 4. Atoms
of one element can combine
with atoms of other elements to form
compounds. A given compound always
has the same relative number and types
of atoms.
 5. Atoms are the units of chemical
change. That is, atoms are not created or
destroyed in chemical reactions. A
chemical reaction simply changes the
way the atoms are grouped together.
 Chemical
Formula: a way of writing a
compound using only symbols and
numbers. The atoms are indicated by he
element symbols and the number of each
type of atom is indicated by a subscript.
 Water- H2O
• Water contains 2 hydrogen atoms and 1 oxygen atom
 Carbon dioxide- CO2
• Carbon dioxide contains 1 carbon atom and 2
oxygen atoms
 The
atom is composed of 3 subatomic
particles called:
 1. Protons
(p+)
 2. Neutrons (n0)
 3. Electrons (e-)
 Protons
and Electrons are always equal in
number in neutral atoms
 Nuclear
Atom: An atom with a dense
center of positive charged around which
tiny electrons move in a space that was
otherwise empty.
 Nucleus: The
relatively small, dense
center of positive charge in an atom.
• Made of Protons and Neutrons
 Proton: A
positively charged subatomic
particle located in the atomic nucleus.
 Neutron: A
subatomic particle with no
charge located in the atomic nucleus.
 Electron: A
negatively charged
subatomic particle located outside of the
nucleus.
Particle
Symbol
Relative
Charge
Relative
Mass
Electron
e-
1-
1 amu
Proton
p+
1+
1836 amu
Neutron
N
0
1839 amu
 Atomic
Number: The number of protons
intside the nucleus of an atom.
• Generally given the symbol Z
 Mass
Number: The sum of the number of
protons and neutrons in the nucleus of an
atom.
• Generally given the symbol A
 The
X
element symbol is given the symbol
A
Z
X
23
=
11
Na
 In
natural samples of an element you may
find atoms with different masses. This
phenomena can be explained by
isotopes
Isotopes: Atoms with the same number of
protons, but with different numbers of neutrons.’
Most atoms have at least two stable isotopes
Exceptions:
Aluminum, Fluorine, and Phosphorus have only 1
Tin has 10
 When
we refer to an isotope we use its
name and mass number
• We don’t have to give the atomic number
because it is the same in all isotopes of a given
element
 Example:
• Boron-10 (10B)
• Boron- 11 (11B)
 Hydrogen
is an exception to the name
and mass number rule
• The isotopes of hydrogen are so important that
they have special symbols and names
 Protium
(P)- hydrogen with no neutrons
 Deuterium (D)- hydrogen with 1 neutron
 Tritium (T)- hydrogen with 2 neutrons
 Percent
abundance of each isotope can
be calculated if the masses of the
isotopes is found using a mass
spectrometer
 Using
a mass spectrometer we find that
the mass of 10B is 10.0129 amu and
11B is 11.0093 amu
 10B
 11B
is 10.0129 amu
is 11.0093 amu
 Average Atomic Mass is 10.811 amu
 By
looking at the information, which isotope
occurs in the greatest abundance?
 10.811=
(%10B •10.0129) + (%11B • 11.0093)
 We
know from algebra that when you add
two percents they must equal 100
• We reduce this to two decimals equal to 1
 % 10B + % 11B = 1, where % 10B = x and % 11B = y
 So, x + y = 1
 We need to have the equation in terms of one variable so,
y=1-x
 10.811=
(x • 10.0129) + ((1 – x) • 11.0093)
 We
find that 10B has an abundance of
19.91% and 11B has an abundance of
80.09%
 This
means that in an average natural
sample of 10,000 boron atoms you would
find that 1,991 would be 10B atoms and
8,009 would be 11B atoms
 Antimony, Sb, has
two stable isotopes
with experimentally determined masses
of 120.904 amu (121Sb) and 122.904 amu
(123Sb). What are the relative abundances
of these isotopes ?
 Periodic
Table: A chart that shows all the
known elements and gives you
information about each one.
• Elements are listed on the periodic table in
order of increasing atomic number
• One of the most useful tools in chemistry
 Elements
are arranged in vertical Groups
and horizontal Periods.
 Periodic
tables used in the United States
have groups numbered 1-8 followed by
the letter: A or B.
A
groups are main group elements
B
groups are transition elements
 Group
1A- Alkali Metals
 Group 2A- Alkaline Earth Metals
 Groups 3B-12B – Transition Metals
 Group 7A- Halogens
 Group 8A- Noble Gases
 The
horizontal periods are numbered
from 1-7
 Period
1 contains only H and He
 Periods 2 & 3 contain 8 elements
 Periods 4 & 5 contain 18 elements
 Periods 6 & 7 contain 32 elements
 The
table is split into 3 basic parts:
 1. Metals
 2. Non-metals
 3. Metalloids/semimetals
 1. Metals:
high electrical conductivity
high luster (shininess)
high ductility (can be drawn into wires)
high malleability (can be rolled into sheets)
can form alloys (solutions of one or more metals
in another metal)
 All
metals are solids except for Mercury
 2. Non-metals:
Nonlustrous
poor conductors of electricity
All lie to the right of the diagonal line that
stretches from B to Te in the periodic table
 Some
are solids; bromine is a liquid, and
a few, like nitrogen and oxygen are
gases at room temp.
 3. Metalloids/ semimetals:
display characteristics of both metals and
nonmetals.
Only silicon, germanium, arsenic, antimony, and
tellurium are in this category
 Most
elements are reactive and are not
found naturally in pure form
 However, there
are a few exceptions
 Gold, Silver, and
Platinum are called
Noble Metals because they are relatively
unreactive
 Group
8 elements are called
Noble Gases because they do not
combine readily with other elements.
• He, Ne, Ar, Kr, Xe, Rn
 Diatomic
Molecules: Molecules made up
of two atoms.
• H2, O2, N2, Cl2, F2, Br2, and I2 are diatomic
molecules in pure, elemental form.
 All
the elements of Group 7 are diatomic
molecules
 Allotropes: Different
forms of elements
where there are the same atoms, but they
are structured differently.
• Carbon comes as Diamond, Graphite, and
Buckminsterfullerene.
 Ion: An
atom or group of atoms that has a
positive or negative charge.
 Taking
a neutral atom and adding or
subtracting one or more electrons can
result in a charged ion.
 Positive
ions are called Cations.
• Produced when an electron is lost from a neutral
atom.
 Mg2+ and Na+ are examples of cations.
 Magnesium normally has 12 protons and 12
electrons, but when 2 electrons are lost it becomes a
cation with a 2+ charge.
 Negative
ions are called Anions.
• Produced when an electron is gained to a neutral
atom.
 Cl- and O2- are examples of anions
 Chlorine normally has 17 protons and 17 electrons,
but when it gains an electron it becomes an anion
with a 1- charge.
 Individual
atoms that have lost or gained
electrons are called Monatomic ions
 Na+
 O2-
11 protons, 10 electrons
8 protons, 10 electrons
 Al3+
13 protons, 10electrons
 Groups
1A- 3A form positive ions with a
charge equal to the group number of the
metal
 1A
Na 1 electron lost Na +
 2A Ca 2 electrons lost Ca2+
 3A Al 3 electrons lost Al3+
 Nonmetals
often form ion with a negative
charge equal to 8-(group #) of the
element.
 5A
N 8-5= 3 electrons gained N3 6A S 8-6= 2 electrons gained S2 7A B 8-7= 1 electron gained
B-
 There
is no easily predictable pattern for
determining charges on transition metals
 Many
of them also form several different
ions
• Iron can be Fe2+ or Fe3+
 Polyatomic
ions contain two or more
atoms with the resulting compound
having an electric charge.
 NH4+
• Four hydrogen atoms surround a nitrogen atom,
and the group has a 1+ charge.
 Many
chemical compounds contain ions.
• We know this because electric currents can run
through them.
• Substances can only conduct electric current if
the ions can move freely
 Salt water has ions, however pure salt and pure water
cannot conduct electricity.
 Ionic
Compounds: A compound that
results when a metal reacts with a
nonmetal to form cations and anions.
• The result must have a net charge of zero.
 1. Both
positive and negative ions must
be present.
 2. The numbers of cations and anions
must be such that the net charge is zero.
 Na+
+ Cl- = NaCl
 Mg2+
 Li+
+ Cl- = MgCl2
+ N3- = Li3N