Lecture 3
The Periodic Table, Atomic Structure,
Isotopes and Ions, Basic Nomenclature : Ch 2
Dr Harris
8/28/12
HW problems: Ch 2: 15, 29, 35, 37, 43, 49, 52, 53
Elements, Molecules and Compounds
• Millions of different materials in the world, all comprised of some
combination of only 118 elements
• Similar to how the alphabet combines 26 letters to yield
hundreds of thousands of words, elements bond in unique
arrangements to give different substances
What exactly is an element?
• Any pure substance that can not be broken down into simpler substances
is an element. Elements combine to form molecules. Molecules
comprise compounds
• Elements: H, O, Na, Cl (etc.)
• Molecules:
• H2O (water),
• HOOH (hydrogen peroxide)
• NaCl (sodium chloride)
• HCl (hydrochloric acid)
• NaClO4 (sodium perchlorate)
• We refer to a compound as a bulk accumulation of molecules (i.e. A
single water molecule vs. a glass full of water composed of billions and
billions of molecules)
Metallic vs Nonmetallic Character
• Elements can be metals, nonmetals, or semiconductors (we will discuss
semiconductors later)
Metal characteristics
•
•
•
•
•
Malleable
Ductile
Conductive of electricity
Conductive of heat
Have luster and shine
Metallic vs Nonmetallic Character
Nonmetal characteristics
• Most nonmetals are gases
• Non conductive
• Those that are solid are not
conductive, malleable, or ductile
Br2 (L)
O2 (g)
P(s)
S(s)
Metals are all elements LEFT of the black line, except Hydrogen.
Elemental Groups and Diatomic Species
• Elements listed down a column have very similar properties, and tend
to behave the same way. These columns are called groups. We will
revisit chemical groups later.
• Certain elements are unstable, and hence, do not commonly exist as
individual species, but as diatomic molecules
• These include H, O, N, and all of the halogens (group 17)
• H  H2 (Hydrogen gas)
• O  O2 (Oxygen gas)
• N  N2 (Nitrogen gas)
• F, Cl, Br, I  F2, Cl2, Br2, I2
(Fluorine gas, chlorine gas, bromine gas and iodine gas)
Solids, Liquids, and Gases
Solids
• Atoms tightly bound
• Fixed volume and shape (does
not conform to container)
• A chemical is denoted as solid by
labeling it with (s)
Au atoms
Au (s)
Solids, Liquids, and Gases
• Liquids
• Atoms less tightly bound than
solids
• Has a definite volume, but not
definite shape (assumes the shape
of its container)
• Denoted by (L) ex. H2O (L)
Water
molecule
Solids, Liquids, and Gases
• Gases
• Free atoms
• No shape, no definite volume
• Can be expanded or compressed
(like engine piston)
• Denoted by (g) ; ex. O2 (g)
O2 molecules
Mixtures
•Most substances in nature are mixtures, chemicals that exist together
without actually bonding. Mixtures can be homogeneous (uniform
throughout) or heterogeneous (non-uniform)
Homogeneous
•
•
•
•
water
fructose
CO2(g)
food coloring
• 78% N2
• 21% O2
Heterogeneous
Separating Mixtures
• In many instances, it is important to separate mixtures. This can be done
by taking advantages of differences in the properties of the components.
Heterogeneous mixture of sugar,
sand and iron. How can we
separate the components??
Critical Thinking
• You’ve received a sample of dirty water from the local pond to purify
in your lab. How could you extract pure water from the contaminated
mix?
You would perform a distillation.
Solutions
• All solutions are homogeneous fluids (free flowing), meaning that
the composition is completely uniform throughout
• Example: salt water. You add salt (NaCl) to water, it completely
dissociates (dissolves; molecules split apart) and spreads evenly
throughout the water. There is the same amount of salt
everywhere in the glass
• Salt is the solute. Water is the solvent.
• When a substance dissolves in water, it forms an aqueous
solution, labeled (aq): Salt water  NaCl(aq)
• Air is another example of a solution
Law of Constant Composition
•
Example
• We analyze 1.630 g of CaS and find that it’s 0.906 g Ca. Find the mass of
S? Find the mass % of Ca and S?
- First, we know that CaS is composed only of Ca and S. Therefore, all of
its mass must come from Ca and S. We can find the mass of S.
•
mass Ca + mass S = total mass of CaS
0.906g + X = 1.630 g
X = 0.724 g S
• Now, we can find the mass % of each element
mass % Ca + mass % S = 100%
Group Examples
• A sample of KI has a mass of 2.00 g and is 23.5% K. What is the mass
% I? What is the mass of I in the compound.
• A sample of NaCl is 39.3% Na, and is found to contain 1.58 g Na.
What is the total mass of NaCl? What is the mass of Cl?
Basic Laws of Chemical Reactions
• Law of conservation of mass: Atoms are not created or destroyed
in a chemical reaction, only rearranged
• Law of constant composition: Specific molecules have fixed
ratios of atoms
-
• One N atom combines with 3 H atoms to make ammonia gas, NH3
(g). It does not matter how much N or H you have.
• So, if you had 10 N atoms and 1000 H atoms, you would make
10 NH3 molecules, with 970 unused H atoms left over.
-Some elements combine in multiple ratios, like C and O
* CO  carbon monoxide
* CO2  carbon dioxide
Dalton’s Atomic Theory
1.
Matter is composed of atoms
2.
The atoms of a given element are exactly the same in every way
3.
The atoms of different elements differ in mass and chemical
behavior
4.
Compounds are composed of two or more atoms of different
elements bonded together.
5.
In a reaction, atoms are rearranged, separated or recombined to
form new substances. No atoms are created or destroyed, and
the atoms themselves DO NOT CHANGE
Atomic Mass Units
• Each element in the periodic table is assigned a relative mass, called the
atomic mass (amu).
• Remember the CaS example. The total mass of CaS is 55.5% Ca and
44.6% S. If all the mass is from Ca and S, then the ratio of the mass %
must also be the ratio of the mass
𝑚𝑎𝑠𝑠 𝐶𝑎 55.6
=
= 1.25
𝑚𝑎𝑠𝑠 𝑆
44.6
• This suggests that a Ca atoms weighs 1.25 times as much as an S atom
• To develop the periodic table, atom masses are determined by using
carbon as a reference
Atomic Mass Units
Calculating Atomic Mass
• What is atomic mass of H2O to 3 sig figs?
• H = 1.0079 amu
• O = 15.9994 amu
2 1.0079 + 15.9994 = 18.0 𝑎𝑚𝑢
• What is the atomic mass of NaHCO3 to 3 sig figs?
• Na = 22.990 amu
• C = 12.01 amu
22.990 𝑎𝑚𝑢 + 1.0079 𝑎𝑚𝑢 + 12.01 𝑎𝑚𝑢 + 3 15.999 𝑎𝑚𝑢 = 68.0 𝑎𝑚𝑢
% Mass Calculations Using Atomic Mass
• What is the % mass of Na in Na2S?
𝑇𝑜𝑡𝑎𝑙 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑁𝑎2 𝑆 = 22.99 𝑎𝑚𝑢 2 + 32.066 𝑎𝑚𝑢 = 78.05 𝑎𝑚𝑢
% 𝑁𝑎 =
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑁𝑎
22.990 𝑎𝑚𝑢 2
=
= 58.90%
𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑁𝑎2 𝑆
78.05 𝑎𝑚𝑢
• How many grams of sulfur atoms are there in a 4.50 g sample of
Na2S?
% 𝑆 = 41.1 %
.411 4.50 𝑔 = 1.85 𝑔 𝑆
Atomic Structure
• Atoms were initially thought to be indivisible, and the smallest
particles in existence
• Toward the end of the 19th century, it was discovered that atoms
consist of even smaller sub-atomic particles
• G. Johnstone Stoney hypothesized that electricity exists in discrete
units, or small packets of equally distributed charge. He called these
electric units electrons
• J.J. Thompson was one of the first scientists to proved the existence
of such particles. He is most famous for his cathode-ray experiment
Cathode Rays
• Electrical discharge (current) in partially evacuated tubes had been
previously observed when high voltages were applied between two
electrodes.
• These rays of current are called cathode rays, because they flowed from
the cathode (negative end) to the anode (positive end)
• Scientist initially thought that these rays were caused by atoms of the
electrodes.
The Electron
• J.J. Thomson modified the typical cathode ray tube experiment in several
ways:
1. He used various materials to compose the cathode. It was found
that the nature of the beams did not depend at all on what the
cathode was composed of, and the mass of the cathode never
changed. Hence, these rays were NOT caused by atoms or heavy
particles
2. He applied either electrical fields (a) or magnetic fields (b) to the
cathode ray. He found that the cathode rays were actually beams of
negatively charge particles that are MUCH lighter than atoms, and
are the same in every element. This marked the discovery of the
electron.
Determination of Electron Mass and
Charge
• Through later experiments, it was determined that the mass of an
electron is 9.109 x 10-31 kg
• The charge of an electron is -1.602 x 10-19 C, where C is the symbol
for coulombs, the SI unit of charge
• For simplicity, we report the relative charge of an electron as -1
Radioactivity
• Shortly after the discovery of electrons, Antoine Henri Becquerel
discovered radioactivity, the process by which atoms spontaneously
break apart.
• By monitoring emissions from a radioactive material under an
electric field, Ernest Rutherford found that radiation consists of
three types of radiation
• alpha (α) particles
• beta (β) particles
• gamma (ϒ) particles
masses below are amu
charges are relative
3 types of Radiation
• Positive α particles drawn
toward negative terminal
• Negative β particles drawn
toward positive terminal
• Neutral ϒ particles are
unaffected by electric field
The “Plum Pudding” Model
• Since atoms are electrically neutral, and
electrons are negatively charged, Thomson
knew that a positively charged sub-atomic
particle must also exist
positive charge
• Thomson envisioned the atom as electrons
evenly distributed in a “sea” of positive
charge, the so-called “plum pudding” model
(shown to the right)
• Ernest Rutherford then formed an
experiment to see how a high energy beam
of positively charged α particles would
interact with atoms
electrons
Discovery of the Nucleus
• A very thin sheet of gold
film was placed in front of a
beam of α-particles
• Most of the α-particles
passed through the gold
film
• However, some α-particles
were scattered at large
angles. How could this
happen?
• Rutherford realized that most of the mass of an atom is concentrated
in a small positively-charged region. He called this region the
nucleus.
Rutherford’s Model of the Atom
Nucleus
Electron
cloud
α particle
• Most of the volume of an atom is empty space, which electrons
spread sparsely throughout. The α-particles easily pass through
this space.
• However, the nucleus is very dense, and very positively charged.
The α-particles that approach this region are strongly repelled.
Protons and Neutrons Comprise the
Nucleus
• Later experiments showed that protons and neutrons reside in the
nucleus and comprise the bulk of the mass of an atom
• Protons and electron have equal but opposite charge
Particle
Relative Charge Mass (amu)
Charges shown in table are relative to the charge
of a proton. A proton has an actual charge of
1.602 x 10-19 C, an electron has a charge of -1.609
x 10-19 C. Opposite charges attract!
Like charges repel!!
Atomic Number
• The number of protons in an atom is
called the atomic number.
• For a given element, the number of
protons DOES NOT CHANGE. As
shown to the right, only Carbon has
6 protons.
• In a neutral atom, the number of
protons is equal to the number of
electrons.
Mass Number
• The mass number of an element is
the sum of the protons and
neutrons.
• The mass numbers listed on the
periodic table are average values.
• The reason for these averages is that
elements exist in nature as multiple
isotopes.
Isotopes
• Isotopes are variations of elements
with the same number of protons
but different numbers of neutrons.
• For example, the most common
isotope of hydrogen contains one
proton, one electron, and no
neutrons (99.985% of all hydrogen
atoms).
• An isotope of hydrogen is
deuterium, which has one neutron
(.0115 %)
• A third isotope, tritium, has two
neutrons (~ 0%).
mass number
1
1𝐻 hydrogen
atomic number
2
1𝐻 deuterium
3
1𝐻 tritium
Average atomic mass is obtained using the %
abundance and the isotope mass.
𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 =
𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑥 (𝑖𝑠𝑜𝑡𝑜𝑝𝑖𝑐 𝑚𝑎𝑠𝑠)
Example
• Revisiting the previous example of the hydrogen
isotopes, we found the following isotopic masses
and abundances in the table.
• Using these values, calculate the average atomic
mass of Hydrogen. Does it match the value given in
the periodic table?
ISOTOPE
% Abundance
Isotopic mass (amu)
1
1𝐻
99.9850
1.00782503
2
1𝐻
0.0115
2.01410178
3
1𝐻
~0
3.01604927
𝑎𝑣𝑔. 𝑚𝑎𝑠𝑠 = .999850 1.00782503 + .000115 2.01410178 + (0)(3.01604927)
= 1.0079 𝑎𝑚𝑢
Group Problem
• The 3 main isotopes of Carbon are C-12, C-13, and C-14. Fill in the
table below with the correct number of protons, neutrons, and
electrons.
• Then, calculate the average atomic mass. Does your value match that
of the periodic table?
ISOTOPE
Protons Neutrons Electrons % Abundance Isotopic mass (amu)
𝟏𝟐
𝟔𝑪
98.93
12
𝟏𝟑
𝟔𝑪
1.07
13.003 354 8378
𝟏𝟒
𝟔𝑪
~0
14.003 2420
How To Determine % abundance
• The average atomic mass of Boron, as given on the periodic table, is
10.811 amu
• It is known that two isotopes of Boron exist,
• 105𝐵
(isotopic mass: 10.013 amu)
•
11
5𝐵
(isotopic mass: 11.009 amu)
• What are the % abundances of each isotope?
Continued.
• We know that:
𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 =
𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑥 (𝑖𝑠𝑜𝑡𝑜𝑝𝑖𝑐 𝑚𝑎𝑠𝑠)
• We also know that the total % abundance must be 100%
• We allow %( 105𝐵) = 𝑥
• Therefore: %( 115𝐵) = 1 − 𝑥
10.811 = 𝑥 10.013 𝑎𝑚𝑢 + (1 − 𝑥)(11.009 𝑎𝑚𝑢)
• Solving for x, we find that
%( 105𝐵) = 19.9 %
%( 115𝐵) = 80.1 %
Ions
• Thus far, we’ve learned than an element is essentially defined by it’s
atomic number
• Each element has an exact number of protons.
• For example, Hydrogen has only one proton. If you force a
second proton onto the atom, you no longer have hydrogen… you
now have Helium.
• We have also learned that atoms of a particular element can have
variations in the number of neutrons. Atoms of the same element
with varying numbers of neutrons are called isotopes.
• Next, we will discuss ions.
Ions
• Ions are electrically charged atoms, resulting from the gain or loss
of electrons.
• Positively charged ions are called cations. You form a cation when
electrons are lost
• Negatively charged ions are called anions. You form anions when
electrons are gained
• An cation is named by adding the word “ion” to the end of the
element name
• Anions are named by adding the suffix –ide to the end of an element
Ions
𝐿𝑖 +
Lithium ion
𝑁𝑎+
Sodium ion
𝑀𝑔2+
Magnesium ion
𝐴𝑙 3+
Aluminum ion
𝐶𝑙 −
Chloride
𝑆 2−
Sulfide
𝑂2−
Oxide
𝑃3−
Phosphide
Ions
32
16𝑆
Gains 2 electrons
Sulfur-32
ISOTOPE
32 2−
16𝑆
Sulfide-32
Protons
Neutrons Electrons
32
16𝑆
16
16
16
32 216𝑆
16
16
18
27
13𝐴𝑙
Loses 3 electrons
Aluminum-27
ISOTOPE
27 3+
13𝐴𝑙
Aluminum ion-27
Protons
Neutrons Electrons
27
13𝐴𝑙
13
14
13
27
3+
13𝐴𝑙
13
14
10
Nomenclature
• There are special rules to naming molecules. In this chapter, you will
see two types of molecules: IONIC and COVALENT
• Ionic bonds are formed between metals and nonmetals
• To name an ionic compound, you do the following
• write the name of the metal
• Follow it with the name of the nonmetal, but change the
ending of the name of the nonmetal to –ide
• Example: KF
• K is a metal. F is a non metal.
• We write the name as: Potassium Fluoride
Na2S = Sodium Sulfide; MgO = Magnesium Oxide
Nomenclature
• Covalent bonds are formed between
nonmetals and nonmetals
• To name a covalent compound, you do
the following
• Include a prefix corresponding to
the first nonmetal. Do not use
mono on the first nonmental.
• Use a prefix to name the second
nonmetal. We only use –mono for
oxygen. Drop the ending of the
second nonmetal and replace it
with –ide
SF4 = sulfur tetrafluoride
N2O = dinitrogen monoxide
P2O3 = diphosphorus trioxide
SiC = silicon carbide
Rules of Hydrogen
• Hydrogen is strange. It’s a nonmetal, but can sometimes act as a metal and is
listed on the metal side of the periodic table.
1. If hydrogen is listed first, and the nonmetal is of group 6, treat it as a metal and
name the compound accordingly (except water).
Ex. H2S = hydrogen sulfide
Compounds of group 17 are strong acids and are named as such. We drop –gen
and end the second non metal with –ic acid
HCl = hydrochloric acid ; HF = hydrofluoric acid
2. If hydrogen is listed last, treat the molecule as covalent. Change the ending to
–ide and use the appropriate prefix.
NaH = sodium hydride; AlH3 = aluminum trihydride
AsH3 = Arsenic trihydride;