Download Atoms: The Building Blocks of Matter

Atoms:
The Building Blocks of Matter
Do Now
• What is a theory?
• What is a model?
• How do you make inferences about things you
can not see?
Atom
• From Greek, meaning
“indivisible”
• Atomic theory (from
400 BCE) = atoms are
the building blocks of
matter
• There wasn’t any
evidence for nearly
2000 years.
Atomic Size
• At sea level, one cubic
centimeter of air (size of a
sugar cube, or marble) will
have 45 billion billion atoms
within it.
– 45,000,000,000,000,000,000
• If you tried to count to
45,000,000,000,000,000,00
0 it would take you 400,000
years
• Fill Rye COMPLETELY with
45,000,000,000,000,000,00
0 marbles.
Atomic Size
• To see the atoms in a drop
of water, you would need
to enlarge the drop until…
it is 24 kilometers wide!
• Think of a line 1 millimeter
long. If this line were
blown up to the size of the
empire state building, an
atom would be…
a tenth the thickness of a
sheet of paper.
Some History
• Democritus
– 460-371 B.C.
– ancient Greek philosopher
– believed all matter consisted of
extremely small particles that
could not be divided
– atoms, from Greek word
atomos, means “uncut” or
“indivisible”
• Aristotle
– believed all matter came from
only four elements—earth, air,
fire and water
Some Scientific Laws
• The Law of Definite Proportions
• The Law of Conservation of Mass
Law of Definite Proportions
• Two samples of a chemical compound contain
the same elements in exactly the same
proportions by mass regardless of the size of
the sample or source of the compound.
– NaCl = 39.3% sodium 60.7% chlorine
– H2O = 11.2% hydrogen 88.8% oxygen
– C2H6O2 = 38.7% carbon, 9.7% hydrogen, 51.6% oxygen
Law of Conservation of Mass
• Mass is neither created nor destroyed during
ordinary chemical reactions or physical
changes
• Thus, the mass of the reactants equals the
mass of the products
Dalton’s Atomic Theory
• Dalton proposed an explanation for the law of
conservation of mass, the law of definite
proportion, and the law of multiple
proportions.
• He reasoned that elements are composed of
one kind of atom and that only whole
numbers of two or more kinds of atoms can
combine to form compounds.
• He proposed the solid sphere model
Dalton’s Atomic Theory
1. All matter is composed of extremely small particles
called atoms, which cannot be subdivided, created, or
destroyed.
2. Atoms of a given element are identical in physical
(size/mass) and chemical properties.
3. Atoms of different elements differ in physical
(size/mass) and chemical properties.
4. Atoms of different elements combine in simple,
whole-number ratios to form chemical compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged, but never created,
destroyed, or changed.
Modern Atomic Theory
• Atoms are divisible into even smaller particles.
These smaller parts of the atom are called
subatomic particles
– Electrons
– Protons
– Neutrons
Discovering Electrons
• In 1897, J.J. Thomson used a cathode ray tube
(passes electricity through a glass tube with
little pressure) to deduce the presence of a
negatively charged particle.
J.J. Thomson’s cathode ray tube
• He knew that rays must have come from the
atoms of the cathode because most of the
atoms in the air had been pumped out of the
tube. Because the cathode ray came from the
negatively charged cathode, Thompson
reasoned that the ray was negatively charged.
J.J. Thomson’s cathode ray tube
• He observed that when a small paddle wheel was placed in
the path of the rays, the wheel would turn. This observation
suggested that the cathode ray consisted on tiny particles that
were hitting the paddles of the wheel.
• His experiments showed the cathode ray consists of particles
that have mass and a negative charge. These were called
electrons.
• He proposed the plum pudding model.
• Cathode Ray Tube
http://www.bbc.co.uk/history/british/victorians/launch_ani_paddle_steamship.shtml
Discovering the Nucleus
• In the 1900s, Ernest Rutherford performed his
gold foil experiments
– He directed small, positively charged alpha
particles (that are helium nuclei) at a thin gold foil.
– Particle hits on the detecting screen (film) were
recorded and deflected angles were measured.
Rutherford’s gold foil experiments
This diagram shows the expected
result of Rutherford's experiment if
the "plum pudding" model of the
atom is correct.
This diagram shows the actual
result. Most of the alpha particles
are only slightly deflected, as
expected, but occasionally one is
deflected back towards the source
Only a very concentrated (dense) positive charge in a tiny space within the gold
atom could possibly repel the fast-moving, positively charged alpha particles
enough to reverse the direction of the alpha particles.
Rutherford’s gold foil experiments
• His experiments showed that the nucleus is very
small and positively charged.
• He also hypothesized that the mass of the
nucleus must be larger than the mass of the
alpha particles, otherwise the alpha particles
would have knocked the nucleus out of the way.
• He also argued that most of the alpha particles
were not deflected, because most of the atom
was empty space.
• He proposed a planetary model or nuclear model
Rutherford’s Gold Foil Experiment
• http://www.mhhe.com/physsci/chemistry/ess
entialchemistry/flash/ruther14.swf
• Rutherford's Gold Foil Experiment
RUTHERFORD ACTIVITY
HALLWAY
Pennies
Rolled Marbles
Rolled Marbles
http://www.learner.org/resources/series61.html
Pennies
Equal
Distance
From each
other
The Nucleus
• Using measurements from
Rutherford’s experiment,
scientists calculated the
radius of the nucleus to be
less than 1/10,000 of the
atom.
– If the nucleus were the size of
a marble, the atom would be
the size of a football stadium
The Nucleus
• Protons = positively charged particles
– The charge of a proton was calculated to be equal but
opposite to the charge of an electron
– The mass of a proton is almost 2000x the mass of an
electron
• Neutrons = neutral particles
– The mass of a neuron is almost equal to the mass of a
proton
• The sum total of masses of protons, neutrons,
and electrons equals the mass of the atom.
Mass of atoms are measured in Atomic Mass Units!
1 amu = 1/12 mass the Carbon-12
(amu)
Atoms
• All living things are made up of tiny units
called ATOMS.
ATOMS consist of electrons orbiting around a
nucleus.
ELECTRONS
• (-) negative electrical charge found in the
space around the nucleus
NUCLEUS
• PROTON
(+) has a positive electrical charge.
• NEUTRON
has a neutral charge (no charge)
Subatomic Particles
ATOM
NUCLEUS
ELECTRONS
PROTONS
NEUTRONS
POSITIVE
CHARGE
NEUTRAL
CHARGE
NEGATIVE CHARGE
Atomic Identity
• If the number of electrons equals the number
of protons, the atom is electrically neutral.
(No electrical charge)
• Elements differ in their number of protons
and therefore in the amount of positive
charge their nuclei possess.
• The number of protons determines an atom’s
identity.
Atomic Number
Atomic # = p+
• Atomic Number (Z)= number of protons of
each atom of that element
– Atoms of different elements have different
numbers of protons (different atomic numbers)
– Atoms of the same element all have the same
number of protons (same atomic numbers)
• The atomic number identifies the element
– 113 elements have been identified, with 113
different atomic numbers
Atomic Number
• Because atoms are neutral, they must have
the equal numbers of protons and electrons.
• Therefore, the atomic number tells us how
many protons and also how many electrons an
atom has.
How many PROTONS
and ELECTRONS are in:
•
•
•
•
•
•
•
Silver
Hydrogen
Neon
Gold
Boron
Sodium
Tungsten
47
1
10
79
5
11
74
Mass Number
Mass # = p+ + n0
• Mass number = the total number of protons and
neutrons (total number of particles in the nucleus)
• Mass numbers can vary among atoms of a single
element, because atoms of the same element can have
different numbers of neutrons.
• Different elements can have the same mass numbers,
because the mass number does not help you identify
the element, the atomic number does!
Try this:
Mass # = p+ + n0
Element
p+
Oxygen
33
Phosphorus
n0
e- Mass #
Nuclear Symbols
•
•
235U
92
235 is the mass number of Uranium
92 is the atomic number of Uranium
• A uranium nucleus has 92 protons.
• It also has a total of 235 neutrons and protons in its
nucleus (mass number).
• How many neutrons in an atom of Uranium-235?
• Mass # – Atomic # = # of Neutrons
• 235 (protons + neutrons) – 92 protons =
143 neutrons
FIND THE NUMBER OF NEUTRONS:
•
•
•
•
•
•
•
Sodium
Calcium
Nitrogen
Iron
Argon
Lithium
What does this tell you?
12
6
C
Modern Atomic Theory
• Atoms of a particular element do share the same
atomic number (number of protons) and identical
chemical properties but the atoms of a given
element may differ in their mass numbers (number
of protons and neutrons).
• Elements occur in nature as mixtures of isotopes.
Isotopes
• Isotopes = atoms of the same element with
different numbers of neutrons and mass
numbers
 Nuclear symbol:
Mass #
Atomic #
12
6
 Hyphen notation: carbon-12
C
Isotopes
• Isotopes = atoms of the same element with
different numbers of neutrons and mass
numbers
 Nuclear symbol:
Mass #
Atomic #
14
6
 Hyphen notation: carbon-14
C
Isotopes
© Addison-Wesley Publishing Company, Inc.
Try to determine information
about these isotopes:
• Chlorine-37
– atomic #:
17
– mass #:
37
– # of protons:
17
– # of electrons:
17
– # of neutrons:
20
37
17
Cl
Isotopes of Hydrogen
Isotope
Hydrogen–1
(protium)
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium)
Protons Electrons Neutrons Nucleus
Using Mass Numbers
• How many protons, neutrons, and electrons make up an atom of Br-80?
•
•
•
•
Protons + Neutrons = 80
Protons = 35
Electrons = 35
Neutrons = 80 – 35 = 45
• How many protons, neutrons, and electrons make up an atom of C-14?
•
•
•
•
Protons + Neutrons = 14
Protons = 6
Electrons = 6
Neutrons = 14 – 6 = 8
Ions
Are created when an atom loses or gains one
or more electrons; it acquires a charge
http://web.visionlearning.com/custom/chemistry/animations/CHE1.3-an-ions.shtml
Charge of Ion = number of protons – number of electrons
More electrons than protons = negative charge (anion)
More protons than electrons = positive charge (cation)
12
6
C +1
# of protons 
# of electrons 
Total charge 
PRACTICE IONS
Ion
Li
+1
Ni
+2
Pb
+2
Ca
+2
Cs
+1
# protons
# neutrons # electrons
Chemical
Symbol
Number of
protons
Number of
electrons
Number of
neutrons
35
36
45
Atom or Ion?
I
11
12
atom
55
78
atom
14
12
Zr
12
50
Br
53
atom
44
atom
Ce
ion
27
25
32
84
80
125
73
68
108
50
71
Sc
Pb
Ni
atom
Do Now
• If your first quarter grade is based 10% on
homework, 20% on labs, 20% quizzes, and
50% tests, what should your grade be if you
averaged 100 on homework, 90 on labs, 80 on
quizzes, and 70 on tests?
–
–
–
–
10% x 100 = 10
20% x 90 = 18
20% x 80 = 16
50% x 72 = 36
– Total = 80 %
Atomic Mass
• Atomic mass = the weighted average of the
masses of the existing isotopes of an element.
• Don’t get these confused!
– Mass number = the total number of protons and
neutrons that make up the nucleus of an atom.
– Atomic mass includes the masses of the protons,
neutrons and electrons of atoms and isotopes.
Weighted Average
• You have a box containing two sizes of
marbles.
• 25% of the marbles have masses of 2.00 g
each
• 75% of the marbles have masses of 3.00 g
each
• Calculate the weighted average….
Calculate the weighted average
• Assume you have 100 marbles
– 25%, or 25, have a mass of 2.00 g
– 75%, or 75, have a mass of 3.00 g
–
–
–
–
25 marbles x 2.00 g = 50 g
75 marbles x 3.00 g = 225 g
Total mass = 50 + 225 = 275 g
275g / 100 marbles = 2.75 g/ marble
– A simpler method is as follows:
25% = 0.25
75% = 0.75
(2.00g x 0.25) + (3.00g x 0.75) = 2.75g
Atomic Mass
• 1 amu = 1/12 the mass of a 12C atom
Carbon = 12.011
Isotope
Symbol
Carbon-12
12C
Carbon-13
13C
Carbon-14
14C
Composition % Abundance in
of the nucleus
nature
6 protons
98.89%
6 neutrons
6 protons
1.11%
7 neutrons
6 protons
<0.01%
8 neutrons
Calculating Atomic Mass
Avg.
Atomic
Mass
(mass)(% )  (mass )(% )

100
• STEP 1: Take the mass # (in amu) of each element
and multiply by its percent abundance (%)
• STEP 2: Add all of these values together
• STEP 3: Divide by 100
Calculating Atomic Mass
• Boron exists as 2 isotopes: B-10 or B-11
• B-10
10
• B-11
11
5B
5B
% Abundance
19.78%
80.22%
Calculating Atomic Mass
Atomic Mass of Boron
• STEP 1: 10 x 19.78 = 197.8
11 x 80.22 = 882.42
• STEP 2: 197.8 + 882.42 = 1080.22
• STEP 3:
1080.22 = 10.802 amu
100
Calculating Atomic Mass
• Calculate the Atomic Mass of Chlorine:
% Abundance
• Chlorine – 35
75.53%
• Chlorine – 37
24.47%
Calculating Atomic Mass
• Calculate the Atomic Mass of Silicon:
% Abundance
• Si – 28
92.21%
• Si – 29
4.70%
• Si – 30
3.09%
Calculating Atomic Mass
• Calculate the Atomic Mass of Oxygen:
% Abundance
• O-16
99.762%
• O-17
0.038%
• O-18
0.200%
Calculating Atomic Mass
• What is the average atomic mass of Cu which
is found in nature as
69.15% Cu-63 (62.929601 amu) and 30.85%
Cu-65 (64.927794 amu) ?
(0.6915 x 62.929601 amu) +
(0.3085 x 64.927794 amu) = 63.55 amu
Relating masses in grams to numbers of atoms
• Mole = the amount of a substance that contains
as many particles as there are atoms in exactly 12
grams of carbon-12.
– The mole is a counting unit, like a dozen.
– The mole relates to masses of atoms and compounds.
• Avogadro’s number = The number of particles in
exactly one mole of a pure substance.
– 6.022 x 1023 particles
• Molar mass = the mass of one mole of a pure
substance; numerically equal to the atomic mass
of the element in atomic mass units (g/mol)
• 1 mole of any element is its atomic mass in grams
What is the molar mass of Li?
6.94 g/mol
What is the molar mass of Hg?
200.59 g/mol
The molar mass of an element
contains one mole of atoms.
4.00g He, 6.94g Li, and 200.59g Hg all
contain one mole of atoms.
How many atoms is this?
Avogadro’s number: 6.02 x 1023 particles
(atoms)
What is the mass in grams of
3.50 mol of Cu?
3.50 mol Cu x 63.55g Cu
1 mol Cu
= 222g Cu
What is the mass in grams of
3.42 mol Ag?
What is the mass in grams of
0.876 mol Pb?
A chemist produced 11.9 g of
Aluminum. How many moles of Al
were produced?
11.9 g Al x 1 mol Al
26.98 g Al
= 0.441 mol Al
How many moles of Na are in
4.01 g Na?
How many moles of Zn are in 0.674 g
Zn?
How many moles of Ag are in
23
3.01 x 10 atoms of Ag?
3.01 x 1023 atoms Ag x 1 mol Ag
6.02 x 1023 atoms Ag
= 0.500 mol Ag
How many atoms are in
32 g of S?
Gram Formula Mass (GFM)
• Gram formula mass is the molar mass, or atomic mass
of a compound
• Units are grams/mole (g/mol)
Example: H2O
1) What is the molar mass of H? O?
• H = 1.007
• O = 15.994
2) Multiply by the subscripts
• For Hydrogen  (1.007) x (2) = 2.014, of 2
• For Oxygen  (15.994) x (1) = 15.994, or 16
3) Add the masses together
• 2+ 16 = 18, the molar mass of H2O
• What is the gram formula mass (GFM) of salt
(NaCl)?
• What is the gram formula mass (GFM) of sugar
(C6H12O6)?
Discoveries about the atom
Dalton
1.
All matter is composed
Of extremely small particles which
cannot be subdivided, created or
destroyed.
2. Atoms of a given element are
identical in physical and chemical
properties.
3. Atoms of different elements have
different physical and chemical
properties.
4. Atoms of different elements
combine in simple whole-number
ratios to form chemical compounds.
5. In chemical reactions, atoms are
combined, separated, or rearranged,
but never created, destroyed, or
changed.
JJ Thomson
What did he
discover: Electron
His Experiment:
Cathode Ray Tube
His findings:
Electrons are
negatively charged
embedded in a
positive charge.
Rutherford
What did he
discover: The Nucleus
His experiment:
GOLD FOIL
EXPERIMENT
(1900’s)
His findings:
The atom is mostly
empty space.
The nucleus is small.
The nucleus is dense.
The nucleus is
positively charged
Niels Bohr
•Electrons revolve around the nucleus
in specific orbits, or energy levels.
• An atom has energy levels. Electrons
can only exist in these energy levels,
not in between.
•When an atom is in the ground state,
the electrons exist in the energy
levels closest to the nucleus.
•GROUND STATE: the lowest energy
state of an atom; the electrons
occupy energy levels closest to the
nucleus.
•If an atom receives, energy, the
atom becomes excited and electrons
jump to higher energy levels.
•EXCITED STATE: an atom with
higher potential energy than in the
ground state because electrons have
“jumped” to a higher energy level.
Solid
Sphere
Model
Electron
Cloud
Model/
Quantum
Model
This model
suggested that
electrons could be
considered waves
confined to the
space around a
nucleus.
Electron cloudsregions where
electrons are
likely to be found
Refining Nuclear Models
• In 1913, Danish physicist, Niels Bohr, refined
Rutherford's idea by adding that the electrons were
in orbits around the nucleus. Rather like planets
orbiting the sun. With each orbit only able to contain
a set number of electrons.
• He proposed a Bohr model or Orbit model
Bohr’s ATOM
Atom
HELIUM
_______
_______
_________
+
_________
N
N
+
-
__________
The Bohr Model
• 1. Electrons revolve around the nucleus in specific
orbits (shells), or energy levels.
• 2. An atom has energy levels. Electrons can only
exist in these energy levels, not in between.
• 3. When an atom is in the ground state, the
electrons exist in the energy levels closest to the
nucleus.
• 4. If an atom receives, energy, the atom becomes
excited and electrons jump to higher energy
levels.
http://www.visionlearning.com/library/flash_viewer.php?oid=1347&mid=51
Ground State
• The lowest energy state
of an atom
• Electrons in the first
energy level have the
lowest potential energy
since they are located
closest to the nucleus.
Excited State
• An atom has a higher potential energy than in the
ground state because electrons have “jumped”
(moved up) to a higher energy level.
• Electrons with higher potential energy occupy orbits
farther from the nucleus. The further an electron is
from the nucleus, the greater its energy!
Modern Atomic Theory
Bohr Model—shows
electrons in orbit
around protons and
neutrons
Quantum-mechanical
model—doesn’t show exact
location of electrons, just
probable place
Current Atomic Model
• Electrons act like particles (because they have a mass) and waves
(because they have certain frequencies corresponding to their
energy levels)
• Electrons are located in orbitals around the nucleus that correspond
to specific energy levels
• Electron clouds = orbitals that do not have sharp boundaries, but
shows 3D region where electrons are most probable to be found.
• Electron Cloud Model or Quantum Model proposed by Louis de
Broglie & Erwin Schrodinger
Electron Configuration
• Arrangement of electrons
• Each atom has a distinct electron
configuration.
• The ground state electron configuration is
found on the periodic table in the lower left
hand corner of each box.
Classify the following as ground state electron
configurations or excited state electron configurations.
Element
ground state electron configurations or
excited state electron configuration
lithium
1-2
calcium
2-8-7-3 excited
chlorine
2-8-7
ground
aluminum
2-7-4
excited
neon
2-7-1
excited
sodium
2-8-1
ground
potassium
2-8-7-2 excited
exited
The configuration listed on the periodic table is the
ground state electron configuration. In the chart below,
draw the ground state and excited state electron
configurations
Element
He
O
Na
F
Al
Mg
Br
Ground State
Excited State
Element
Ground State
Electron
Configuration
Ion
Ion’s Electron
Configuration
Na
Na
+
Mg
Mg
+2
Fe
Fe
+3
Al
Al
+3
Li
Li
+1
Quantum
• Electrons can only absorb or release energy in
discrete, specific amounts.
• The amounts, or bundles of energy are called
quanta (or photons) corresponding to
differences in energy levels of the
orbitals/shells.
• The greater the radius of an orbit (the farther
from the nucleus), the greater the energy of
the electrons in that orbit. The orbits or shells
are known as principal energy levels.
Electrons and Light
• An atom emits energy when the electron falls from
high energy levels to lower energy levels. This energy
is in the form of electromagnetic radiation.
• If the wavelength is in the visible light spectrum, the
energy can be seen as color.
Light Emission
• Each move from a particular energy level to a
lower energy level will release light of a
specific wavelength.
• When certain elements are excited, they give
off energy of a distinctive color as the electron
fall back down to lower energy levels. These
colors are specific and can be used to identify
the elements (Flame Test).
Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum
…produces a “bright line” spectrum
Spectral lines
• If high voltage is applied to hydrogen gas confined in a gas
tube, called a gas discharge tube, light is emitted. If this light
is passed through a prism, a series of bright lines of distinct
colors is produced. Bohr reasoned that these different colored
bands of light were actually quanta of corresponding energy.
These quanta were emitted as electrons of hydrogen atoms
returned from their higher levels in the excited state to their
lower levels in the ground state.
Bright Line Spectra
• Bright line spectrum =
the series of bright lines
produced when excited
electrons return to their
original energy levels
• Each element has its
own unique set of
spectral lines which can
therefore be used to
identify the elements
presence.
Valence Electrons
Electrons that occupy the valence energy level
Valence Electrons= found in outer most energy level
Na 2-8-1
Cl 2-8-7
Atoms can have a maximum of 8 valence electrons
Lewis Dot Diagrams
(Electron Dot Diagrams)
• Represent the arrangement of
electrons around the nucleus.
• Electrons are the DOTS.
• Nucleus is the symbol.
• ONLY REPRESENT VALENCE
ELECTRONS!!
• Fill one side first, then one on each
side before you pair electrons.
Lewis Dot Diagrams
(Electron Dot Diagrams)
Na
B
O
Mg
Cl
Ne
Si
H
N