Download Chapter 2a - Angelfire

Unit 2 – Chapters 2, 8, & 9
The Components of Matter
Definitions for Components of Matter
________________ - the simplest type of substance with unique
physical and chemical properties. An element consists of only
one type of atom. It cannot be broken down into any simpler
substances by physical or chemical means.
__________________- a structure that
consists of two or more atoms that are
chemically bound together and thus behaves
as an independent unit.
Definitions for Components of Matter
__________________ - a substance
composed of two or more elements
which are chemically combined.
________________ - a group of
two or more elements and/or
compounds that are physically
intermingled.
• __________________
(1766-1844), an
English schoolteacher
and chemist, studied
the results of
experiments by
Lavoisier, Proust, and
many other scientists.
Dalton’s Atomic Theory
• Dalton proposed his atomic theory of matter
in 1803.
• Although his theory has been modified
slightly to accommodate new discoveries,
Dalton’s theory was so insightful that it has
remained essentially intact up to the present
time.
Dalton’s Atomic Theory
The Postulates
1. All matter consists of ______________________.
2. Atoms of one element ___________________
be converted into atoms of another element.
3. Atoms of an element are _____________________
in mass and other properties and are
_______________________ from atoms of any
other element.
4. ______________________________ result from the
chemical combination of a specific ratio of atoms of
different elements.
Structure of the Atom
• J.J. (John Joseph) Thomson, physicist
• 1890-1900
• Showed that the atoms of any element can be
made to emit tiny negative particles - called
_______________________.
• Thompson knew that the entire atom was not
negatively charged so he concluded that the atom
must also contain positive particles that balance
the negative charge, giving the atom a
___________________________________.
• Ernest Rutherford
• 1911
• Learned physics in J.J. Thomson’s
laboratory in the late 1890s.
• Main area of interest was the
_____________________________ positively charged particles with a
mass approximately 7500 times that
of an electron.
Ernest Rutherford
• By 1919, Rutherford concluded that the
nucleus of an atom contained what he
called ______________________(has the
same magnitude of charge as the
electron, but its charge is positive)
• Protons have a _________ charge and the
electron a charge of ______________.
• 1932, he and a coworker (James
Chadwick) were able to show that most
nuclei also contain a neutral particle that
they named the _________________
(which has no charge)
Modern Concept of Atomic Structure
• The simplest view of the atom is that it
consists of a tiny nucleus that is about 10-13
cm in diameter.
• Electrons move about the nucleus at an
average distance of about 10-8 cm from it.
• Nucleus contains ___________________,
which have a positive charge equal in
magnitude to the electron’s negative
charge, and ________________________,
which have almost the same mass as a
proton but no charge.
Modern Concept of Atomic Structure
• Mass and charge of the electron (e-),
proton (p+), and neutron (N)
The mass and charge of the electron, proton, and neutron.
Particle
Relative Mass*
Relative Charge
Electron
1
1Proton
1836
1+
Neutron
1839
None
*The electron is assigned a mass of 1 for comparison
Distinguishing Between Atoms
• Protons and electrons are _______________
in an atom of an element (neutral charge).
• The ________________________ of an
element is the number of ______________ in
the nucleus of an atom of that element. (If the
p+ and e- are the same, then the atomic
number will also identify the number of e-)
Distinguishing Between Atoms
• The sum of the number of neutrons and the number
of protons in a given nucleus is called the atom’s
____________________________.
protons + neutrons = mass number
• _____________________
• atoms with the same number of protons but different
numbers of neutrons.
• Elements on the periodic table are the most common
isotopes of those substances.
Distinguishing Between Atoms
• Isotopes
• Because they have different numbers of
_________________________, their mass
numbers will be different.
• Neon - 20
• Neon - 21
• Neon - 22
• All of these are isotopes of neon.
Distinguishing Between Atoms
• Isotopes
• 3 known isotopes of hydrogen
• hydrogen - 1
• hydrogen - 2
• hydrogen - 3
[hydrogen]
[deuterium]
[tritium]
Isotopic Symbols
• X = the symbol of the element
• A = the mass number
• Z = the atomic number
A
Z
1H
1
2H
1
3H
1
Hydrogen
Deuterium
Tritium
X
Atomic Masses
• Because atoms are so tiny, the normal units of mass the gram and the kilogram - are much too large to be
convenient.
• Mass of a single carbon atom is 1.00 x 10-23 grams.
• When describing the mass of an atom, scientists
have defined a much smaller unit of mass called the
__________________________.
Atomic Masses
• In terms of grams:
• 1 amu = atomic weight of a substance
expressed in grams
• 1 carbon atom = 12.01 amu = 12.01 grams
• 1 aluminum atom = 26.98 amu = 26.98
grams
Periodic Table of Elements
Periodic Table of Elements
• Shows all the known elements and
gives a lot of information about each
element.
• Invaluable in chemistry!
Development of Periodic Table
• Elements in the
same group
generally have
similar
________________
________________.
• Properties are not
identical, however.
Development of Periodic Table
Dmitri Mendeleev
and Lothar Meyer
independently
came to the same
conclusion about
how elements
should be
grouped.
Development of Periodic Table
Mendeleev, for instance, predicted the
discovery of germanium as an element with
an atomic weight between that of zinc and
arsenic, but with chemical properties similar
to those of silicon.
Dmitri Mendeleev
• Organized the elements according to
their increasing
__________________________.
• Then he grouped them into columns
and rows according to physical and
chemical properties.
• Row – __________________
• Column - __________________
Henry Moseley
• Rearranged the elements according to
their __________________________.
• Arranging the elements in this manner
provided for a better fit of chemical and
physical properties and aligned those
elements that were discovered after
Mendeleev developed the original
periodic table.
Parts of the Periodic Table of Elements
• _____________________ – substances
to the left of the dark line
• _____________________ – substances
to the right of the dark line
• _____________________ – those
elements that border the line
Properties of Metal, Nonmetals,
and Metalloids
Metals versus Nonmetals
Differences between metals and nonmetals
tend to revolve around these properties.
Metals versus Nonmetals
• Metals tend to form ____________.
• Nonmetals tend to form _____________.
Metals
Tend to be
______________,
______________,
______________,
and good
conductors of
______________
and
______________.
Metals
• Compounds formed
between metals and
nonmetals tend to
be
______________.
• Metal oxides tend to
be
______________.
Nonmetals
• ______________,
______________subs
tances that are
______________
conductors of heat
and electricity.
• Tend to gain
______________ in
reactions with metals
to acquire noble gas
configuration.
Nonmetals
• Substances
containing only
nonmetals are
______________
compounds.
• Most nonmetal
oxides are
______________.
Metalloids
• Have some
characteristics of
______________,
some of
______________.
• For instance, silicon
looks shiny, but is
brittle and fairly poor
conductor.
Fireworks
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
Potassium – combustible element that helps oxidize firework
mixtures
Lithium – adds red color
Sodium – gold and yellow colors
Magnesium – bright white color
Calcium – deepens the colors of the other elements in the
fireworks
Strontium – red color and stablizes other elements
Barium – green color and stablizes other elements
Titanium – produces the spark
Iron – produces sparks
Copper – blue color
Zinc – smoke clouds
Aluminum – silver and white sparks and flames – sparklers
Carbon – black powder
Phosphorus – fuel
Sulfur – fuel
Antimony – glitter effects
Groups of the Periodic Table
•
•
•
•
•
Group 1 – Alkali metals
Group 2 – Alkaline Earth metals
Group 11 – Coinage metals
Group 17 – Halogens
Group 18 – Noble Gases
Electronic Structure
Electronic Structure
• ______________ are vital in the
determining the properties of elements.
• Electrons are involved in bonding
between ______________.
• Electrons move around the nucleus in
different
____________________________.
Energy Levels
• There are ______________ energy
levels – one for each corresponding row
or period on the periodic table.
• Within each energy level there are
______________.
Sublevels
• Within each energy level there are a
possibility of 4 sublevels – depending
on which energy level you are dealing
with.
Sublevels
• The 4 different sublevels are:
•
•
•
•
_____ – holds a maximum of 2 electrons
_____ – holds a maximum of 6 electrons
_____ – holds a maximum of 10 electrons
_____ – holds a maximum of 14 electrons
Sublevels
• Energy level 1 – only has an “s” sublevel
• Energy level 2 – only has an “s” and “p”
sublevel
• Energy level 3 – only has an “s”, “p”, and “d”
sublevel
• Energy level 4 & 5 – have an “s”, “p”, “d”, and
“f” sublevels
• Energy level 6 – only has an “s”, “p”, and “d”
sublevel
• Energy level 7 – only has an “s” and “p”
sublevel
Sublevels
•
•
•
•
•
•
Level 1 – maximum of ______________
Level 2 – maximum of ______________
Level 3 – maximum of ______________
Level 4 & 5 – maximum of ___________
Level 6 – maximum of ______________
Level 7 – maximum of ______________
Orbitals
• Within each sublevel, there are orbitals
– locations where the electrons are
actually located.
• An orbital can hold 2 electrons only.
•
•
•
•
Sublevel “s” – _____orbital (total of 2 e-)
Sublevel “p” – _____orbitals (total of 6 e-)
Sublevel “d” – _____orbitals (total of 10 e-)
Sublevel “f” – _____orbitals (total of 14 e-)
Electron Configuration
• Shows the location of the electrons in
an atom of an element.
• ______________– electrons fill specific
energy levels and sublevels from the
nucleus out in a specific order.
Electron Configuration
• 1. Determine the number of electrons
an atom of an element has.
• 2. Fill the energy levels and sublevels
in order using the diagonal rule.
• 3. Write the configuration
Electron Configuration
• Sodium – 11 electrons
• 1s2 2s2 2p6 3s1
• Follow the diagonal for filling the energy
levels and sublevels
• If the sublevel is filled, include all the
electrons, however, when only one is needed
(as in the 3s1), only include one electron.
• If you total the superscript numbers (number
of electrons), there are only 11, which is the
number of electrons that sodium has.
Electron Configuration
• Bromine – 35 e-
• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
• Notice that after the 4s it is 3d. You
must go in the specific order as shown
by the diagonal rule.
Orbital Notation
• ______________shows the electrons within
each energy level and sublevel and how they
______________.
• We know that 2 ______________ can exist
within an orbital.
• We also know that electrons are all
______________ charged.
• For the 2 electrons to exist in an orbital they
must spin in opposite directions.
Orbital Notation
• ______________– electrons will remain
as unpaired as possible.
• Minimizes
____________________________–
everyone gets their own room whenever
possible.
• Pauli’s Exclusion Principle
• Electrons will spin in opposite directions
e- Configuration & Orbital Notation
__ __ __ __ __ __
• 1s2 2s2
2p6
3s1
• The lines above show the maximum
number of orbitals for each sublevel.
• All possible orbitals must be shown
even if electrons do not exist in the
orbital.
e- Configuration & Orbital Notation

__
1s2

  

__
__ __ __ __
2s2
2p6
3s1
• Notice that the arrows go in opposite
directions – shows the electrons spinning in
opposite directions.
• The last energy level / sublevel only has one
arrow because there is only one electron in
that sublevel.
• http://intro.chem.okstate.edu/APnew/Default.html
• Website with simulation
Valence Electrons
• ______________ electrons are the
electrons in the highest energy level, “s”
and “p” sublevel.
• Totaling 8 electrons – 2 in the “s” and 6 in
the “p”
• These are the electrons involved in
______________.
Quantum Numbers
The Uncertainty Principle
• Heisenberg showed that the more precisely
the momentum of a particle is known, the
less precisely is its position known.
• In many cases, our uncertainty of the
whereabouts of an electron is greater than
the size of the atom itself!
• We can only calculate the probability of
finding an electron within a given space.
Quantum Mechanics
• Erwin Schrödinger
developed a
mathematical
treatment into which
both the wave and
particle nature of
matter could be
incorporated.
• It is known as
_________________.
Quantum Numbers
• Each orbital describes a spatial
distribution of electron density.
• An orbital is described by a set of three
___________________.
Principal Quantum Number, n
• The principal quantum number, n, describes
the ____________________ on which the
orbital resides.
• The values of n are integers of numbers
ranging from 0 to infinity.
• The larger the n value, the larger the
______________.
• As n increases, the energy increases.
• Lower energy = more __________ atom (n
would be small)
Principal Quantum Number, n
• An electron can jump from a lower
energy state to another by emitting or
absorbing ______________.
• Absorb energy – electrons can move to a
higher energy state.
Azimuthal Quantum Number, l
• This quantum number defines the
_______________ of the orbital.
• We use letter designations to
communicate the different values of l
and, therefore, the shapes and types
of orbitals.
• 0 to n-1
Azimuthal Quantum Number, l
Value of l
0
1
2
3
Type of orbital
s
p
d
f
Magnetic Quantum Number, ml
• Describes the three-dimensional
orientation of the orbital.
• Values are integers ranging from -l to l
• Assign the “blanks” in orbital notation
with zero on the middle blank and then
–l through zero to +l.
Magnetic Quantum Number, ml
• Orbitals with the same value of n form a shell.
• Different orbital types within a shell are
subshells.
s Orbitals
• Value of l = 0.
• Spherical in shape.
• Radius of sphere
increases with
increasing value of n.
p Orbitals
• Value of l = 1.
• Have two lobes with a node between them.
d Orbitals
• Value of l is 2.
• Four of the
five orbitals
have 4 lobes;
the other
resembles a
p orbital with
a doughnut
around the
center.
Spin Quantum Number, ms
• In the 1920s, it was
discovered that two
electrons in the same
orbital do not have
exactly the same
energy.
• The “spin” of an
electron describes its
_________________,
which affects its
energy.
Spin Quantum Number, ms
• This led to a fourth
quantum number,
the spin quantum
number, ms.
• The spin quantum
number has only 2
allowed values: +½
and -½.
Spin Quantum Number, ms
• Electrons spin on an ____________
• Electron spin quantum number
• Electrons in the same orbital spin in
opposite directions
• Values of +½ and –½
Pauli Exclusion Principle
• No two electrons in
the same atom can
have exactly the
same energy.
• For example, no two
electrons in the same
atom can have
identical sets of
quantum numbers.
• The first electron placed in an orbital
gets the +½ and the second one gets
the -½.
Sulfur
• Electron Configuration for sulfur is:
• 1s2 2s2p6 3s2p4
• Highest energy level is 3, so n = 3
• The last electron is in the p orbital, so l = 2
 

• ___
___ ___
-1
0
+1
• The last electron is in the -1 position, so ml = 1
• That electron is the 2nd electron in that orbital,
so it has an ms = - ½
Quantum Numbers for Sulfur
• 3, 2, -1, - ½