Chemistry 100
Chapter 2
Atoms, Molecules and Ions
Law of Conservation of Matter


By 1800, chemists had noted that
the mass of reactants equals the
mass of products - provided you
capture any escaping gas
Matter is not created or destroyed
in a chemical reaction
Law of Constant Composition

Chemists (notably Proust)


The relative amount (percentage) of
each element in a compound was the
same no matter how the compound
was made
These two laws lead Dalton to
revive the Atomic Theory

Matter is made up of small, indivisible
particles
Dalton’s Atomic Theory


An element is composed of atoms.
All atoms of a given element are the
same.

Atoms of different elements are different
and have different properties.
Dalton’s Theory-II


Atoms are not changed, created or
destroyed in a chemical reaction.
Compounds are the combination of
more than one element.

A given compound has the same relative
number and kind of atoms.
Law of Multiple Proportions

Accorrding to Dalton’s Theory




Two elements (A and B) form two distinct
compounds
The amounts of B combining with a fixed amount
of A would be a small whole number ratio.
Water: 1 g hydrogen + 8 g oxygen
Hydrogen peroxide: 1 g hydrogen +16 g
oxygen
Atomic Structure
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
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Roentgen discovered X-rays (1895)
Becquerel discovered radioactivity
(1896)
J. J. Thompson discovered the
electron (1897)
Rutherford classified radioactivity
emissions: alpha (), beta () and
gamma ()
Alpha, beta, gamma



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Rutherford’s findings:
Alpha are positive particles (+2) ;
heavier than electrons
Beta are high speed electrons;
negatively charged particles (-1)
Gamma are neutral rays
Alpha particles are nucleus of He
atom
Thompson’s Model
“Plum pudding” model
A cloud of positive charge holding the
negatively charge electrons in place
Rutherford’s alpha experiment



Scattering of alpha particles by gold
foil
Most particles were undeflected
Some were deflected by large angles
Rutherford’s Explanation
Rutherford’s Model



Centre - the nucleus - is small but
positively charged
Most of the atom is empty spaces
Electrons rotate about the nucleus like the solar system
Modern Model

Additional experiments showed



Nucleus consists of protons (positive)
and neutrons (neutral)
Electrons (negative) exist around the
nucleus
Number of protons = number of
electrons
Mass of Elementary Particles


Protons and neutrons have a mass of
about 1 amu
Electrons have very small mass

Most of the mass of an atoms comes
from nucleus (1 amu is 1.66054  10-24
grams)
Quantum Mechanical Model
Atomic Mass & Atomic Number
A is Mass Number,
protons + neutrons
A
Sy
Z
Z is Atomic Number,
number of protons
Isotopes




All atoms of a given element have the same
number of protons. All Carbon atoms have
6 protons (and 6 electrons).
The number of protons is different for each
element.
Atoms of a given element that differ in the
number of neutrons are called isotopes.
Examples : carbon-12 and carbon-14
12
6
C
14
6
C
Atomic Masses


By international agreement, the
carbon-12 atom is defined as
having a mass of exactly 12 atomic
mass units (amu’s).
All atomic masses are referenced to
this standard.
The Periodic Table

A typical entry in the periodic table
Atomic number
20
Ca
40.078
Atomic mass
Periodic Table (II)

Elements in the periodic table are
arranged in




Groups or families – they have similar
chemical and physical properties
Metals – towards the left
Nonmetals – towards the right
Metalloids – in the middle region
Atomic Masses in the Periodic Table


Question: why is the mass of C in
the periodic table reported as 12.01
amu and not as 12.000 … amu,
exactly?
Another example: the atomic mass
of Cl is = 35.453 amu’s. We would
expect Cl to be  35 amu?
Ionic Compounds


Tables of common ions in textbook
(pages 60 and 63).
Ionic compounds


Cation name followed by anion name,
e.g., sodium bromide (NaBr)
Multiple ion types


FeCl2 – iron (II) chloride
FeCl3 – iron (III) chloride
Binary Molecular Compounds

Binary compounds containing two
nonmetals


name of the first element in the formula
followed by the stem of the name of the
second element with the suffix -ide.
The number of atoms of each element in
the compound is indicated by a prefix.
mono
di
tri
tetra
1
2
3
4
Some common names must be
committed to memory.
Examples – nitrous oxide, N2O,
and nitric oxide, NO.
Formulas and Names of Acids



An acid usually is a compound of
hydrogen and a nonmetal or a polyatomic
anion.
Treat the hydrogen atoms of the acid as
H+ ions.
For acids containing monatomic anions,


When these acids are found in water solution,
add the prefix hydro- and the suffix -ic to the
stem of the name of the anion
Hydrofluoric acid (HF), hydrochloric acid (HCl)
Acids From Polyatomic Anions

If the anion name ends in ‘ate’, the ‘ate’ in
the name of the anion is replaced by ‘ic
acid ’



The acid of the sulfate ion is sulfuric acid (H2SO4)
The acid of the nitrate ion is nitric acid (HNO3)
If the anion names end in ‘ite’, we change
the suffix to –ous and add the word acid.


The acid of the nitrite is called nitrous acid, HNO2
The acid of the hypochlorite ion is called
hypochlorous acid, HClO
Average Atomic Masses


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Most elements in nature exist as mixtures
of isotopes.
Atomic masses reported in the periodic
table - weighted averages of the different
isotopes.
The amount of each isotope in a sample of
as an element may vary considerably with
the source of the sample.

This is the reason why some elements in the
periodic table have few significant figures for
their mass.
Organic Compounds

Many organic compounds have
complex three dimensional
structures



chains, and/or rings
branches.
The highlighted groups are called
functional groups.

They are primarily responsible for the
chemical and some physical properties
of the molecules.
Alkanes
Methane (CH4)
H
C
H
Ethane (C2H6)
Propane (C3H8)
H
H
H
H
C
C
H
H
H
H
H
C
Butane (C4H10)
H
H
C
C H
H H
H
H
H
H
H
H
C H
C
C
C
H
H
H H H
Alcohols
Methanol (CH3OH)
Ethanol
(CH3CH2OH)
1-Propanol
(CH3CH2CH2OH)
2-propanol
(CH3CH(OH)CH3)
H
H
H
C OH
H
H H
C C
H H
OH
H H
C C C OH
H
H H H
H
OH
H
C
C H
H C
H
H
H
H
Organic Amines
Methyl Amine (CH3NH2)
H
H
Dimethyl amine
((CH3)2NH)
Ethyl Amine
(CH3CH2NH2)
C NH2
H
CH3
HN
CH3
H H
H C
C
H H
NH2
Organic Acids
Formic (Methanoic) acid
(HCOOH)
Acetic (Ethanoic) Acid
(H3COOH)
Propionic (Propanoic) Acid
(H3CCH2COOH)
O
C OH
H
H O
H
C C
H
OH
H O
C C C OH
H
H H
H