Download ATOMS: THE BUILDING BLOCKS OF MATTER

ATOMS: THE BUILDING BLOCKS OF MATTER
EARLY HISTORY OF CHEMISTRY
Leucippus, Democritus
World is made of ________________________________________ (Greek – Atomos;
indivisible)
Aristotle
Matter was ________________________________
Matter does not consist of smaller particles
Before 16th century
Alchemy:
Attempts (scientific or otherwise) to __________________________
17th century
First “chemist” to perform quantitative experiments
Published “The Skeptical Chymist”
Definition of elements became generally accepted
18th century
________________ Phlogiston flows out of a burning material
________________ Discovered oxygen and was found to support combustion
Antoine Lavoisier
Law of Conservation of Mass: ___________________________________________
Combustion involves oxygen, not “phlogiston”
Joseph Proust
Law of Definite Proportion: _____________________________________________
___________________________________
______________________
Invention of a simple set of symbols for the elements, along with system of writing
formulas and compounds
Discovered elements Cerium, Thorium, Selenium, and Silicon
John Dalton (1808)
Proposed an explanation for the law of conservation of mass, law of definite
proportion and law of multiple proportion
Postulates of Dalton’s Theory
All matter is
Atoms
Atoms of a given element are
Atoms of different elements combine in
In chemical reactions, atoms are
LAW OF MULTIPLE PROPORTION
Atoms of different elements combine in simple whole number ratios to form chemical
compounds
Not all aspects of Dalton’s atomic theory have proven to be correct
Dalton’s atomic theory have been modified
All matter is composed of atoms, which are ____________________________________
A given element can
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STRUCTURE OF THE ATOM
J. J. Thomson (Cathode Ray Research 1897)
Cathode rays consist of small negatively charged particles called electrons
J. J. Thomson (Cathode Ray Research 1897)
Electrons are like raisins dispersed in a pudding (the positive charge cloud)
Robert Millikan
(Oil Drop Experiment, 1909)
Determined the charge of the electron and confirmed that the electron carry
negative charge
Calculated mass of electron: 1/1840 mass of hydrogen atom (9.11 x 10-31 kg)
GOLD FOIL EXPERIMENT (1911) – By Ernest Rutherford, Hans Geiger, & Ernest Marsden
A thin sheet of gold foil was bombarded with
Most of the alpha particles passed straight through
Many particles were
GOLD FOIL EXPERIMENT – Conclusions
Most of the alpha particles passed directly through the foil because the atom is
The deflected alpha particles are those that had a
NUCLEAR ATOM MODEL
An atom with a
THE NUCLEUS
PROTONS
Positively charged subatomic particle
Discovered a beam consisting of positive particles called “protons” using a
modified cathode ray tube
J. J. Thomson
Protons have the same amount of electrical charge as an electron but opposite
in charge (+1)
Calculated the mass of proton – 1840 times that of electron
Mass of proton : 1.673 x 10 –27 kg
NEUTRONS
Subatomic particles with no charge
James Chadwick (1932)
Uncharged particles with a mass slightly greater than protons were formed
when Beryllium was bombarded with high-energy alpha particles
Named it neutron
Mass: 1.675 x 10-27 kg
NUCLEAR FORCE
Short-range proton-neutron, proton-proton, proton-neutron force that holds the
nuclear particles together
Nucleus has a high density – 2 x 108 metric tons/cm3
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NUCLEUS
Located near the center of the atom
Has a positive charge
Occupies a very small part of the atom
Has very high density
Composed of protons and neutrons
ELECTRON CLOUD
Region occupied by electrons
Cloud of negative charge
Surrounds the nucleus of the atom
Occupy most of the volume of the atom
OTHER SUBATOMIC PARTICLES
LEPTON (light)
HADRONS (heavy)
Baryons –
Mesons – quarks, antiquarks
Quarks: Up, down, charm, strange, bottom, top
Quarks are held together by gluons
Murray Gell-Mann discovered quarks
ANTIPARTICLES
EX: electron (-) and positron (+)
SIZES OF ATOMS
Atomic Radius
Distance from the
Atomic radii range from
FORCES IN ATOMS
GRAVITY
Carried
by
Graviton
WEAK
w + w - z0
ELECTROMAGNETIC
Photon
STRONG
Gluon
Acts on
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COUNTING ATOMS
Atomic Number (Z)
Number of
Determines the identity of the element
Henry Moseley (1913)
Studied X-rays produced in X-ray tubes with
X-ray wavelength depend on
Determined the
Mass Number (A)
Total number of protons and neutrons in the nucleus of the atom (Rounded-off atomic
mass)
Atomic Mass - mass of protons, neutrons and electrons in a single atom (when the atom is
motionless)
ISOTOPES
Atoms of the same element that have the
Usually identified by specifying the mass number (hyphen notation)
Carbon-12, Carbon-14,
Uranium-235, Uranium-238
Sodium-23, Sodium-24
Hydrogen-1, Hydrogen-2, Hydrogen-3
Nuclide
General term for
Nuclear Symbol
Shows the composition of the nucleus
EXAMPLES:
PROTONS, NEUTRONS, & ELECTRONS
Atomic # =
Mass # =
# Neutrons =
Identify elements using atomic #
Mass numbers of elements CAN change but Atomic numbers will always be the same
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Protons, Neutrons, & Electrons
Nuclide
Symbol
Atomic
Mass
# of p+
number number
# of n0
# of e-
Aluminum-27
108
47
25
9
4
Iodine-131
Be
9
30
4
53
RELATIVE ATOMIC MASSES
Carbon-12 Nuclide
Standard used for units of atomic mass
Atomic Mass Unit (amu)
1/12 the mass of a carbon-12 atom
Mass Spectrometer
Instrument used to determine the relative atomic mass of atoms by the deflection of
their ions on a magnetic field
Average Atomic Mass
Weighted average reflects both the ___________ and the _______________________
of the isotopes as they occur in nature
AveAtomicM ass 
(% abundance A  AtMass A )  (% Abundance B  AtMass B )
100
63
65
Cu
Cu
29
The element copper is found to contain the naturally occurring isotopes
and 29 . The
relative abundances are 69.1% and 30.9% respectively. Calculate the average atomic mass of
copper.
36
38
40
Ar
Ar
Ar
Three isotopes of argon occur in nature – 18 , 18 , and 18 . Calculate the average
atomic mass of argon to two decimal places given the following relative atomic masses of
each of the isotopes” argon-36 (35.97 amu; 0.337%), argon-38 (37.96 amu; 0.063%), and
argon-40 (39.96 amu; 99.600%).
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Gallium has two naturally occurring isotopes, Ga-69 and Ga-71, with masses of 68.9257 amu
and 70.9249 amu. Calculate the percent abundances of these isotopes of gallium.
MASS, MOLE & MOLAR MASS
Mole (mol)
Amount of substance that contains as many particles as there are in 12 g of carbon-12
Avogadro’s Number (6.022 x 1023units)
Number of particles in exactly one mole of a pure substance
Units can be
Molar Mass (g/mol)
Mass in
It is equal to the atomic mass of the element
Examples
Carbon – 12.01 g/mol (1mol C=12.01 g)
Neon – 20.18 g/mol (1 mol Ne = 20.18 g)
GRAM/MOLE CONVERSIONS
1 mol = Molar Mass in grams (Atomic Mass)
1 mol = 6.022 x 1023 atoms
MOL TO GRAM
What is the mass in grams of 3.50 mol of the element Silicon, Si?
GRAM TO MOL
A chemist produced 15.0 g of aluminum, Al. How many moles of aluminum were produced?
ATOMS TO MOL
How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver?
MOL TO ATOMS
How many atoms of Nickel, Ni, are in 5.00 mol of Ni?
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ATOMS TO MASS
What is the mass in grams of 7.50 x 1015 atoms of Zinc, Zn?
MASS TO ATOMS
How many atoms of sulfur, S, are in 12.5 g of sulfur?
THE PERIODIC TABLE
GROUPS OR FAMILIES
Elements in the same group have the same properties
PERIODS OR SERIES
TYPES OF ELEMENTS
_____________ – elements that are hard and shiny, good conductors of heat and electricity
_____________ – elements that are non-lustrous, brittle, and poor conductors of heat and
electricity
_____________ – elements that border the zigzag line in the periodic table
Have some characteristics of both metals and nonmetals
_____________ – elements of Group 18 which are generally unreactive
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