Unit 2 Lecture

Report
Chapter 2
Atoms, Molecules, and Ions
Chemistry, The Central Science, 13th edition
Theodore L. Brown, H. Eugene LeMay, Jr., Bruce E. Bursten,
Murphy, Woodward, and Stoltzfus.
2.1 The Atomic Theory
of Matter
Greek Philosophers

Early philosophers debated the
fundamental “stuff” of which the
universe was made

Democritus’ idea of the world was
based on “atomos” meaning
indivisible or uncuttable

Later Plato and Aristotle
formulated that there can be no
uncuttable “things” and they
successfully propagated their
theory and therefor became the
dominant theory for centuries
John Dalton

The atom reemerged in the 17th
century

The ground work was laid by
John Dalton when he noticed that
elements that reacted with other
elements to form new compounds

Dalton’s atomic theory was based
on 4 postulates
Dalton’s Atomic Theory

Explains several laws that were known during Dalton’s time, including the
law of constant composition.


The Law of Conservation of Mass


In a given compound, the relative numbers and kinds of atoms are constant
The total mass of materials present after a chemical reaction is the same as the
total mass present before the reaction
The Law of Multiple Proportions

If two elements A and B combine to form more than one compound, the masses of
B that can combine with given mass of A are in the ratio of small whole numbers
Dalton’s Atomic Theory
Dalton’s Theory Correct?
Give it some thought

Compound A contains 1.333 g of oxygen per gram of carbon, whereas
compound B contains 2.666 g of oxygen per gram of carbon.

What chemical law does this data illustrate?


Law of Multiple Proportions
If compound A has an equal number of oxygen and carbon atoms, what can we
conclude about the composition of compound B?

There are twice as many oxygen atoms in compound B than there is in compound A
2.2 The Discovery of the
Atomic Structure
Subatomic Particles

Dalton based his theory on observations made in the laboratory


As time progressed, scientists began to probe the nature of matter and
started to discover subatomic particles


He and those who followed had no direct evidence of the atom
We will see that the atom is composed in part by electrically charged particles
Keep in mind as we continue that same charges repel one another, whereas
particles with unlike charges attract one another
Cathode Rays and Electrons (e-)

During the 1800’s scientists experimented (including Thomson) with
evacuated glass tubes with electrodes inserted at both ends

Once a charge was applied a radiation between the electrodes was produced

Cathode rays emanated from the negative end and traveled to the positive
end

The rays were unseen but caused certain materials to fluoresce

The rays were tested with magnets and electrically charged rods and it was found
that the cathode rays were negatively charged

The identity of the new particle was the same regardless of the gas used


This new “thing” was called the electron
Thomson tested the beam and found that there was 1.76 x 108 coulombs per gram
Give it some thought

Thomson observed that the cathode rays produced in the cathode–ray tube
behaved identically, regardless of the particular metal used as cathode. What
is significance of this observation?

All atoms have the same subatomic particles called electrons!
Millikan Oil Drop Experiment

Once the charge-to-mass ratio of
the electron was known, Milikan
was then able to experimentally
figure out the mass of the electron

Using an experiment similar to the
one pictured to the right he
calculated the mass of the electron
to be 9.10 x 10-28 g by solving the
charge of a single electron

This showed that the electron has
a mass of about 2000 times less
than hydrogen!
Radioactivity



In 1896, Henri Becquerel discovered that uranium emitted radiation.

Spontaneous emission is known as radioactivity

He concluded that the source or the radiation was the uranium atoms
Studies done by Ernest Rutherford showed that there were 3 types of
radiation

Alpha (α)

Beta (β)

Gamma (γ)
Alpha and beta radiation was shown to be bent by an electric field (but in
different directions) while gamma radiation was unaffected by it
α and β Rays

Both considered fast moving particles

Beta rays were shown to be the radioactive equivalent of cathode rays



Attracted to positively charged plates!

Charge of -1
The alpha particles were shown to have a positive charge

Charge of +2

Has a mass 7400 times that of an electron
Gamma radiation is high-energy radiation similar to X-rays and does not
consist of particles and carries no charge
3 Types of Radiation
Thomson Model

Thomson in the early 1900’s
reasoned that since the electrons
are such a small portion of the
mass of the atom that it must also
only make up a small portion of the
atom

Thomson proposed that that the
electrons were embedded in the
atom like raisons in pudding

The “pudding” having a positive
charge
Rutherford Atomic Model

Rutherford was studying the angles at
which alpha particles were being
deflected

He discovered that most particles passed
straight through the gold foil while only a
small amount of the particles were
deflected or even bounced back

Rutherford explained the result by
postulating the nuclear model of the
atom, which most of the mass of each
gold atom resided in the nucleus

He postulated that the majority of the
space in an atom was empty and that
the charge of the nucleus was positive
Protons and Neutrons

Subsequent experiments led to the discovery of positive particles (protons)
and neutral particles (neutrons) in the nucleus.

Protons were discovered in 1919 by Rutherford and neutrons in 1932 by British
scientist James Chadwick (1891–1972). Thus, the atom is composed of
electrons, protons, and neutrons.
Give it Some Thought

What happens to most of the alpha particles that strike the gold foil in
Rutherford’s experiment?


They passed straight through
Why do they behave that way?

The atom is mostly empty space
2.3 The Modern View of
Atomic Structure
The Modern View of Atomic Structure

The charge of an electron is -1.602 * 10-19 C.

The charge of a proton is opposite in sign but equal in magnitude to that of an
electron: +1.602 * 10-19 C.

The quantity 1.602 * 10-19 C is called the electronic charge.


For convenience, the charges of atomic and subatomic particles are usually
expressed as multiples of this charge rather than in coulombs.

The charge of an electron is 1- and that of a proton is 1+.

Neutrons are electrically neutral
Every atom has an equal number of electrons and protons, so atoms have no
net electrical charge.
Atomic Diameter

Protons and neutrons reside in the
center of the atom and take up
only a small volume of the atom

Most atoms have diameters
between 1 * 10-10 m 1100 pm2 and
5 * 10-10 m 1500 pm2. A
convenient non–SI unit of length
used for atomic dimensions is the
angstrom 1A° 2, where 1 A° = 1 *
10-10 m. Thus, atoms have
diameters of approximately 1 - 5
A° . The diameter of a chlorine
atom, for example, is 200 pm, or
2.0 A° .
Give it Some Thought

If an atom has 15 protons, how many electrons does it have?


Phosphorous
Where do the protons reside in an atom?

The nucleus
Subatomic Particles
Sample Exercise 2.1

The diameter of a U.S. dime is 17.9 mm, and the diameter of a silver atom is
2.88 A° . How many silver atoms could be arranged side by side across the
diameter of a dime?
Atomic Numbers, Mass Numbers, and
Isotopes

The atoms of each element have a
characteristic number of protons.
The number of protons in an atom
of any particular element is called
that element’s atomic number.


Because an atom has no net
electrical charge, the number of
electrons it contains must equal
the number of protons.
The atomic number is indicated by
the subscript; the superscript,
called, the mass number is the
number of protons and neutrons
Isotopes

Isotopes differ by the number of neutrons but contain the same number of
protons
2.4 Atomic Weights
Atomics Mass Scale

Scientists in the 19th century were aware that atoms of different elements
have different masses


They knew that in water that there was 88.9 g of oxygen for every 11.1 grams of
hydrogen in 100 grams of water
Today we can determine the masses of individual atoms with a high degree of
accuracy.

The atomic mass unit is a convenient way when dealing with very small masses

i.e 1 mole of carbon twelve weights 12 amu
Atomic Weight

Atomic weight is the natural mixture of isotopes found in nature

To calculate the atomic weight of an element is to use the following equation:

Atomic Weight = Σ[(isotope mass) x (fractional isotope abundance)]

Σ sigma means over all masses
Give it Some Thought

A particular atom of chromium has a mass of 52.94 amu, whereas the atomic
weight of chromium is given as 51.99 amu. Explain the difference in the two
masses.
Mass Spectrometers
Some Pictures From Mr. Hunter’s
Graduate Research
Actual Data Gathered by Mr. Hunter
Why are there bands over a certain mass range and not just a line?
The mass spectrometer is an energy filter that allows only stable trajectories of ions through.
Therefore, all ions of the same mass have a range of energies (different velocity) coming from
the ion source.
Book Work in Class

Pages 74-75

Answer only red questions for section 2.1-2.4
2.5 The Periodic Table
The Periodic Table

As the list of elements began to grow throughout history, certain elements
were noticed to have similar traits as other elements.

For example, Na and K had similar properties while He and Ne also had similar
properties to one another.

These similarities became known as periodic trends or families
Periodic Table
Periods= horizontal rows
Families (groups) = vertical columns
Common Families
Metals vs Nonmetals
Metals are found to the left of the stair-step line
Non-metals are found to the right of the stair-step line
Metalloids are found along the stair-step line
2.6 Molecules and
Molecular Compounds
Molecules and Molecular Compounds
All of the pictured molecules are
in their chemical formulas except
for one
Oxygen and hydrogen exist as diatomic elements,
i.e molecule
Molecules that contain more than one type of atom
are called molecular compounds
Molecular and Empirical Formulas

Chemical formulas that indicated the exact number of atoms in a molecule
are called molecular formulas

Chemical formulas that indicate the relative number of atoms in a molecule
are called empirical formulas
Picturing Molecules

Structural formulas shows which
atoms are attached to which

Atoms can be represented by their
symbols and bonds by lines

This type of modeling does not show
the actual geometry of the molecule

The Ball-and Stick model has the
advantage of showing the actual
geometry of the atoms in relation to
one another

The space filling model shows the
relative sizes of atoms and their
geometries
Book Work

Page 75-76

Answer only red questions for section 2.5-2.6
2.7 Ions and Ionic
Compounds
Ions and Ionic Compounds



An ion is a charged particle

A cation is a positive particle (one or more less
electrons)

An anion is a negative particle (one or more
extra electrons)
Ions form because certain elements have a
weak attraction to the outer electrons (valence
electrons)

Other elements have a strong attraction for their
electrons and surrounding elements electrons

The stronger pull can take the weaker element’s
electron

This transfer of electrons causes ions to form
In addition to simple ions there are polyatomic
ions such as ammonium and sulfate

These ions are made of at least 2 elements
Predicting Ionic Charges

Elements like to have a stable number of electrons.

The most stable elements on the periodic table are the noble gases because
they have a full octet of electrons

Elements gain or lose electrons to get to this stable state


Nonmetals tend to gain electrons

Metals tend to lose electrons
Half Filled shells are more stable that partially filled shells

Meaning, electrons can move from one shell to another to stabilize the ion

This also means that s-orbital electrons can be taken instead of f- or d-shell
electrons
Practice Problem

Predict the charge states of:

Fe

Sn

Sc

Ti

V
Ionic Compounds
Ionic compounds are a compounds made up of anions and cations
Ionic compounds are generally metals and nonmetals.
Criss-Cross Method

Since there are not discrete ionic “molecules” we are only able to write
empirical formulas for these compounds

The criss-cross method take the charges of each ion and puts it down at the
subscript of the opposite ion
2.8 Naming Inorganic Compounds
Naming Inorganic Compounds

Chemical Nomenclature comes from the latin word nomen (name) and calare
(to call)

The rules of chemical nomenclature are based on the division of substances
into categories

Organic compounds contain carbon and hydrogen often with oxygen, nitrogen, and
other elements

All other compounds are inorganic compounds
Names and Formulas of Ionic Compounds
Most metals that form cations can form different oxidation states. These are usually transition metals
Oxyanions- an anion containing
one or more atoms bonded
to another element (as in the
sulfate and carbonate ions)
Oxyacids- an acid “see above“
Names and Formulas of Acids
Names and Formulas of Binary Molecular
Compounds

The name of the element farther to the left in the periodic table (closest to
the metals) is usually written first. An exception occurs when the compound
contains oxygen and chlorine, bromine, or iodine (any halogen except
fluorine), in which case oxygen is written last.

If both elements are in the same group, the one closer to the bottom of the
table is

named first.

The name of the second element is given an -ide ending.

Greek prefixes indicate the number of atoms of each element.

(Exception: The prefix mono- is never used with the first element.) When the
prefix ends in a or o and the name of the second element begins with a vowel, the
a or o of the prefix is often dropped.
Give it Some thought

Is SOCl2 a binary compound?
Book Problems

Answer on red question for section 2.7-2.8
2.9 Some Simple Organic Compounds
Organic Chemistry

The study of compounds containing
carbon is called organic chemistry

Organic chemistry is a very
important branch of chemistry
encompassing a large portion or
what chemists do

Being able to name these
compounds is very important for
the standardization of chemistry
experiments

Hydrocarbons are composed of only
carbon and hydrogen

Alkanes are the simplest of
hydrocarbons

Alkanes always end in -ane
Some Derivatives of Alkanes

Functional groups can be added to
organic molecules to change their
chemical properties

Alcohols are made by adding an
–OH group in place of a hydrogen

An alcohol has the ending –ol

Numbers are used to determine
what carbon the functional group
is attached

Always use numbers that make it
the smallest possible number or
smallest possible addend
Isomers

Isomers are compounds that have
the same molecular formulas but a
different chemical arrangement

The most common type of isomer
are structural isomers
Other Functional Groups

There are hundreds of different
functional groups in organic
chemistry

Common functional groups are:

Halogens

Hydroxyl (-OH)

Carboxylic acid (-COOH)

Oxygen insertions

Carbonyl groups (-CO)

Alkane groups
Common Functional Groups
Give it some thought

Draw the structural formulas of two isomers of butane, C4H10

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