Transcript Football ra region center

11th grade
Physical Science
Chapter 4 Notes
Section 1:
Section 2:
Section 3:
Section 4:
Atomic Structure
A Guided Tour of the Periodic Table
Families of Elements
Using Moles to Count Atoms
SECTION 1: ATOMIC STRUCTURE
Objectives:
1. Explain Dalton’s atomic theory and describe why it was more
successful than Democritus’s theory.
2.
State the charge, mass, and location of each part of an atom
according to the modern model of the atom.
3.
Compare and contrast Bohr’s model with the modern model
of the atom.
What are Atoms?
1. Atoms are tiny units that determine the properties of all
matter.
ex: Aluminum cans are lightweight and easy to crush
because of the properties of the atoms that make up the
aluminum.
2. It has taken many centuries for us to understand atoms.
a. In 4th century B.C. the Greek philosopher Democritus
suggested that the universe was made of invisible units called
atoms.
b. The word “atom” is derived from the Greek meaning
“unable to be divided”.
c. Democritus believed movements of atoms caused the
changes in matter that he observed.
d. Democritus explained some observations, he was unable to
provide the evidence needed to convince people that atoms
really existed.
Democritus
1. This was Democritus’ atomic model. It was simply a round
sphere with no electrons, protons, or neutrons.
2. Democritus created the first atomic model. His contribution
helped people with understanding the idea of an atom, and
helped other scientists further look into the science of the
atom and its genetic makeup.
This is Democritus' atomic theory exactly:
1. All matter consists of invisible particles called atoms.
2. Atoms are indestructible.
3. Atoms are solid but invisible.
4. Atoms are homogenous.
5. Atoms differ in size, shape, mass, position, and arrangement.
a. Solids are made of small, pointy atoms.
b. Liquids are made of large, round atoms.
c. Oils are made of very fine, small atoms that can easily slip
past each other.
Dalton’s atomic theory:
1. His atomic theory said that elements consisted of tiny
particles called atoms. It states an element is one of a kind (aka
pure) because all atoms of an element are identical.
2. All the atoms that make up the element have the same mass.
3. All elements are different from each other due to differing
masses.
John Dalton’s atomic theory
1. In 1808, English school teacher John Dalton proposed his
own atomic theory.
a. It was developed with a scientific basis, and some parts
of his theory still hold true today.
b. Like Democritus, Dalton proposed that atoms could not
be divided.
c. According to Dalton, all atoms of a given element were
exactly alike. Dalton also stated that atoms of different
elements could join to form compounds.
2. Dalton’s theory is considered the foundation for the
modern atomic theory.
Atoms are the building blocks of molecules
1. An atom is the smallest part of an element that still has the
element’s properties.
a. If you were to break down a copper coin into the smallest
pieces that were too small for you to see and keep dividing it,
you would be left with the simplest units of the coin – copper
atoms.
2. Atoms are made of protons, neutrons, and electrons.
a. Nucleus – center of each atom with a positive electric
charge, made of protons and neutrons.
b. Protons and neutrons are almost identical in size & mass, but
protons are positively charged and neutrons are neutrally
charged (or have no electric charge at all).
c. Electrons – move around outside the nucleus as a cloud of
tiny negatively charged subatomic particles
Size of atoms
*If you enlarged an Atom to the size of a football stadium,
then the nucleus would be the size of a small marble!
Unreacted atoms have no overall charge
1. Atoms have no charge even though they are made of
electrically charged protons and electrons.
2. Atoms do not have a charge because they have an equal
number of protons and electrons (charges cancel out).
3. A helium atoms has 2 protons and 2 electrons. The atom is
neutral because the positive charge of the 2 protons exactly
cancels the negative charge of the 2 electrons.
Charge of 2 protons :
Charge of 2 electrons:
Charge of 2 neutrons:
Overall charge
+2
-2
0
0
Subatomic particles
Protons
Neutrons
Electrons
Models of the atom
1. Like most scientific models and theories,
the model of the atom has been revised
many times to explain such new discoveries.
2. Bohr’s model compares to planets:
a. 1913 – Danish scientist Niels Bohr suggested that electrons in
atoms move in set paths around the nucleus (much like planets
orbit around the sun).
b. In Bohr’s model, each electron has a certain energy that is
determined by its path around the nucleus; this path defines the
electron’s energy level.
c. Electrons can only be in certain energy levels. They must gain
energy to move to a higher energy level, or lose energy to move
to a lower energy level.
Electrons act more like waves
1. By 1925, Bohr’s model of the atom no longer explained
electron behavior.
2. A new model was proposed that no longer assumed that
electrons orbited the nucleus along definite paths like planets
orbiting the sun.
3. In this modern model of the atom, it is believed that
electrons behave more like waves on a vibrating string than like
particles.
An electron’s exact location cannot be determined
1. It is impossible to determine both the exact location of an
electron in an atom and the electron’s speed and direction.
(Compare it to the moving blades of a fan, and determining
where any one of the blades of the fan was located at a certain
instant.)
2. The best scientists can do is calculate the chance of finding an
electron in a certain place within an atom.
a. One way to visually show the likelihood of finding an
electron in a given location is by shading.
b. The darker the shading, the better the chance of finding an
element at that location.
c. The whole shaded region is called an electron cloud.
Electrons exist in energy levels
1. The number of filled energy levels an atom has depends on
the number of electrons.
2. 1st level holds 2 electrons, 2nd level holds 8 electrons, and the
3rd level holds 18 electrons.
3. Ex: Li – has 3 electrons, so it will have 2 in the 1st level and 1
in the 2nd level
4. Sketch the following on your paper:
Electrons are found in orbitals within energy levels
Electrons are found in orbitals within energy levels
1. The regions in an atom where electrons are likely to be found
are called orbitals.
2. Within each energy level, electrons occupy orbitals that have
the lowest energy.
3. The simplest kind is the s orbital.
a. It can have only 1 possible orientation in space because it
is in a sphere shape.
b. It has the lowest energy and can hold 2 electrons.
4. A p orbital is dumbbell shaped and can be oriented 3 different
ways in space.
(Imagine the y-axis being flat on the page. Imagine the dotted
lines on the x- and z-axes going into the page, and the darker lines
coming out of the page.)
a. A p orbital has more energy than an s orbital.
b. Because each p orbital can hold 2 electrons, the 3 p orbitals
can hold a total of 6 electrons.
p
5. The d (below on left) and f (below on right) orbitals are much
more complex.
a. There are 5 possible d orbitals and 7 possible f orbitals.
b. An f orbital has the greatest energy.
6. Even though the orbitals are very different in shape, each can
hold a maximum of 2 electrons.
We categorize electrons according to what orbital level in which they reside.
The four orbitals are s, p, d, and f.
They are classified by divisions on the periodic table, as follows:
Every atom has between 1 & 8 valence electrons
1. An electron in the outermost energy level of an atom is called
a valence electron.
2. Valence electrons determine an atom’s chemical properties
and its ability to form bonds.
3. The single electron of a hydrogen atom is a valence electron
because it is the only electron the atom has.
Ex: In a neon atom, which has 10 electrons, 2 fill the lowest
energy level. Its valence electrons that are farther away from
the nucleus in the atom’s second (and outermost) energy level.
QUIZ!
1. What part of an atom contains the protons & neutrons?
nucleus
2. What are the regions of an atom where electrons are found?
orbitals
3. Who proposed the first widely accepted version of the
atomic theory?
Dalton
4. Who proposed a model in which electrons can only be in
certain energy levels?
Bohr
5. What are electrons in the outermost energy level called?
valence electrons
SECTION 2:
A GUIDED TOUR OF THE PERIODIC TABLE
Objectives:
1. Relate the organization of the periodic table to the
arrangement of electrons within an atom.
2. Explain why some atoms gain or lose electrons to form ions.
3. Determine how many protons, neutrons, and electrons an
atom has, given its symbol, atomic number, and mass number.
4. Describe how the abundance of isotopes affects an
elements average atomic mass.
You will be creating this!
How the Periodic Table is organized
1. The order the elements are arranged on the periodic table is
by the number of protons in their nucleus (atomic number).
a. E.g.: A Hydrogen atom has 1 p, so H is the 1st element listed.
b. A He atom has 2 p and is the 2nd element listed.
2. Elements are listed in this order in the periodic table because
the periodic law states that when elements are arranged this
way, similarities in their properties will occur in a regular
pattern.
The periodic table helps determine
electron arrangement
1. Horizontal rows are called periods.
2. Just as the number of protons an atom has increases by 1 as
you move from left to right across a period, so does its number
of electrons.
a. H and He are both located in Period 1.
b. The first 2 energy levels contain electrons in the s orbital.
3. Li is in period 2, so it has a 3rd ein an s orbital in the 2nd energy level.
Li:
Energy level:
1
2
Orbital:
s
s
# e-:
2
1
4. As you continue to move right in Period 2, a C atom has electrons
in s orbitals and p orbitals.
C atom’s 6 electrons location are as follows:
Energy level
Orbital
# of e- Sketch of C atom:
1
s
2
2
s
2
2
s
2
5. Each orbital can hold 2 electrons.
Label on the sides
(not in the squares)
where each orbital
matches the table
on your blank table.
Elements in the same group have similar properties
1. Valence electrons determine the chemical properties of
atoms.
2. Atoms of elements in the same group, or column have the
same number of valence electrons, these elements have similar
properties.
Label yours
with 1a – 8a
Some Atoms form ions
3. Atoms of Group 1 elements are reactive because their outermost
energy levels contain only one electron.
4. Atoms that do not have their filled outer s and p orbitals may
undergo a process called ionization.
a. Meaning – they may lose or gain valence electrons so that they
have a full outermost s and/or p orbital.
b. If an atom gains or loses electrons, it no longer has the same
number of electrons as it does protons; and it forms an ion with a net
electric charge.
5. A sodium atom loses 1 electron to form a 1+ charged ion.
a. A sodium atom (Na) shown below loses 1 e- to form a sodium ion
(Na+) or cation.
b. A sodium ion (Na+) is
much less reactive than a
Na atom because it has a full
outer orbital.
6. A fluorine atom gains 1 electron to form a 1- charged ion.
a. F is located in Group 17 of the periodic table, and each atom
has 9 electrons (2 in 1st energy level, 7 in the 2nd level).
b. A F atom needs only 1 more electron to have a full outermost
energy level.
7. An atom of fluorine easily gains this electron to form a negative
ion, or anion.
8. Ions of fluorine are called fluoride ions and are written as F-.
9. Because atoms of other Group 17 elements also have 7 valence
electrons, they are also reactive and behave similarly to fluorine.
How do the Structures of Atoms Differ?
1. Because atoms have different structures, they have different
properties.
2. Atomic number is the number of protons in the nucleus (and it
never changes).
3. Electrons and protons are equal numbers in an atom.
4. The mass number equals the total number of subatomic
particles in the nucleus (protons plus neutrons) because protons
and neutrons provide most of the atom’s mass.
5. Isotopes of an element have different numbers of neutrons,
but same number of protons and electrons.
a. Many elements have only 1 stable form, while others have
different “versions” of the same atom.
b. Those different versions, or isotopes, vary in mass but are all
atoms of the same element.
Some isotopes are more common than others
1. Hydrogen is present on both the sun and on Earth.
2. In both places, protium (the hydrogen isotope without
neutrons in its nucleus) is found most often.
3. Only a very small fraction of the less common isotope of
hydrogen, deuterium, is found on the sun and on Earth.
4. Tritium is an unstable isotope that decays over time, so it is
found least often.
Calculating the number of neutrons in an atom
1. If you know the mass number and atomic number, then you
can calculate the number of neutrons an atom has.
2. Uranium has several isotopes. The isotope uses in nuclear
reactors is uranium – 235.
3. Like all U, it has a mass number of 92, so it must have 92 p+
and 92 e-.
a. Its mass number is 235, so its number of p+ and n equals
235.
Mass number
235
Atomic number
-92
_____
Number of neutrons 143
The mass of an atom
1. The mass of a single atom is very small.
2. Because it is very hard to work with such tiny masses, atomic
masses are usually expressed in atomic mass units.
3. An atomic mass unit (amu) = 1/12th of the mass of a carbon12 atom.
a. This isotope of carbon has 6 p+ and 6n, so individual
protons and neutrons must each have a mass of about 1.0amu
because electrons contribute very little mass.
b. Often the average atomic mass is listed for the element on a
periodic table – which is a weighted average for an element of
the element as it is found in nature.
Ch. 4, Section 3:
Families of Elements
Objectives:
1. Locate alkali metals, alkaline-earth metals, and transition
metals in the periodic table.
2. Locate semiconductors, halogens, and noble gases in the
periodic table.
3. Relate an element’s chemical properties to the electron
arrangement of its atoms.
How are Elements classified?
1. Elements are classified as metals or nonmetals.
2. This classification groups elements that have similar physical and
chemical properties.
3. Most elements fall under 3 groups: metals, nonmetals, and
semiconductors.
a. metals – shiny solids that can be stretched & shaped, and good
conductors of heat and electricity
b. nonmetals – all can be found (except H) on the right side of the
periodic table.
(1) they may be solids, liquids, or gases. (2) solid nonmetals are
typically dull, brittle, and poor conductors of heat & electricity
c. Semiconductors – elements that are considered to be their own
group;
(1) sometimes referred to as metalloids; (2) can conduct electricity
under certain conditions
As you color and label your
table – make a key!
Metals
1. The alkali metals (Group 1) are located on the
left edge of the periodic table.
2. They are soft and shiny, and reacts violently
with water.
3. An atom of an alkali metal is very reactive
because it has 1 valence electron that can easily
be removed to form a positive ion.
4. Because alkali metals such as sodium are so
reactive, they are not found in nature as
elements. Instead, they combine with other
elements to form compounds.
e.g.: The salt you use to season your food is
actually the compound sodium chloride, NaCl.
5. Alkali metals can explode if they are exposed
to water so they are stored in oil to escape from
oxygen and water.
Alkaline-earth Metals
1. Group 2 of the periodic table are called alkalineearth metals.
2. They have 2 valence electrons.
3. Alkaline-earth metals are less reactive than alkali
metals, but they may still react to form positive ions
with a 2+ charge.
4. Examples of properties of these metals:
a. Calcium – make up hard shells of many sea
animals, limestone, and human bones and teeth
b. Magnesium – lightest of all structural metals and
used to build some airplanes; activates many of the
of the enzymes that speed up processes in the
human body; and commonly used in medicines –
milk of magnesia and Epsom salt.
Transition Metals
1. Groups 3 – 12 .
2. Transition metals, like Gold, are much less reactive than sodium or
calcium, but they can lose electrons to form positive ions.
3. They all conduct heat & electricity.
4. Most can be stretched & shaped into flat sheets, or pulled into
wire.
5. Examples of common and/or useful transition metals:
a. copper – used for electrical wiring or plumbing
b. tungsten – used in light bulb filaments
c. Fe, Co, Mn all play
vital roles in your body
chemistry.
6. Hg is the only liquid
at room temperature.
**They are in blue on the
table to the right: **
Synthetic elements, Other Metals, and
Lanthanide & Actinide Series (Rare earth)
1. Technetium and promethium are both man-made elements.
They are also both radioactive, which means the nuclei of their
atoms are continually decaying to produce different elements.
2. The last 2 periods of transition metals (Lanthanide & Actinide
series) are placed toward the bottom to keep the periodic table
narrow so that similar elements elsewhere in the table still line up.
3. All elements with atomic numbers greater than 92 are also manmade and are similar to technetium and promethium.
Examples of common uses of these elements:
a. Promethium-147 is an ingredient in some “glow in the dark”
paints.
b. Americium, is radioactive. Tiny amounts of americium-241 are
found in most household smoke detectors; although it contains
small amounts of radioactive material can affect you, it is safe
when contained inside your smoke detector.
“Just a little FYI...”
Sometimes Lanthanide and Actinide series are referred to as rare
earth metals or “other metals”
What element is most metallic?
Which element is
least metallic?
Nonmetals
1. Except for H, nonmetals are found on the right side of the
periodic table.
2. They include some elements in groups 13 – 16, and all the
elements in groups 17 & 18.
3. There are specific names for the elements in group 17 and 18.
Carbon
1. Carbon and other nonmetals are found on the right side of the
periodic table.
2. Carbon is found in 3 different forms and can also form many
compounds.
3. Although carbon in its pure state is usually found as graphite
(“pencil lead”) or diamond, the existence of fullerenes, a 3rd form,
was confirmed in 1990.
a. The most famous fullerene consists of a cluster of 60 C atoms.
b. The way C atoms are connected in this resembles a pattern of
a soccer ball.
4. Carbon is found in both living and
nonliving things (Glucose, chlorophyll,
isooctane in gasoline, and rubber tires)
Halogens
1. Group 17 elements are halogens, which means “salt forming”.
2. All but At are nonmetals, and all share similar properties.
3. A halogen atom typically gains or shares 1 e when it reacts with
other elements.
4. All are very reactive, and the uncombined elements are
dangerous to humans, but quite useful.
a. F – reacts with almost every other known substance,& small
amounts are added to our water to help prevent tooth decay
b. Cl gas (commonly in the form of the compound hypochlorite) is
used in small amounts to kill bacteria in most swimming pools
c. Elemental chlorine is a poisonous yellowish green gas made of
pairs of joined chlorine atoms (Cl2)
d. Compounds of C & F make up the nonstick coating on cookware.
e. Adding a compound containing Iodine as the negative ion
iodide, I- , to table salt makes “iodized” salt, which you need in
your diet for your thyroid gland to function properly.
Noble Gases
1. Are found in group 18 of the periodic table.
2. They are different from most elements that are gases
because they exist as single atoms instead of as molecules.
3. Like other member of Group 18, neon is inert, or unreactive,
because its s and p orbitals are full of electrons.
a. For this reason, neon and other noble gases do not gain or
lose electrons to form ions.
b. They also don’t join with other atoms to form compounds
under normal conditions.
4. Neon, Helium and argon are other common noble gases.
a. He – used in balloons and to lift blimps
b. Ar – used to fill light bulbs because of its lack of reactivity
prevents filaments from burning
c. Ne – responsible for the bright reddish orange light of “neon”
signs.
Semiconductors (Metalloids)
1. Only 6 elements – Boron, Silicon, Germanium, Arsenic, Antimony,
and Tellurium – are semiconductors/metalloids.
2. They are classified as nonmetals, but each one has some properties of
metals.
a. As their name, semiconductors, implies, they are able to conduct
heat and electricity under certain conditions.
3. Silicon is the most familiar semiconductor; it accounts for 28% of the
mass of Earth’s crust.
a. Sand is the most common silicon compound, called silicon dioxide.
b. Silicon is also an important component of other semiconductors
devices such as transistors, LED display screens, and solar cells.
Your table should look similar & Make a key!
Ch. 4 Section 4:
Using Moles to Count Atoms
Objectives:
1. Explain the relationship between a mole of a substance and
Avogadro’s constant.
2. Find the molar mass of an element by suing the periodic table.
3. Solve problems converting the amount of an element in moles
to its mass in grams, and vice versa.
The mole is useful for counting small particles
1. Because chemists often deal with large numbers of small particles,
they use a large counting unit – the mole, abbreviated mol.
2. A mole is a collection of a very large number of particles.
3. It is about 602, 213, 670, 000, 000, 000, 000, 000!
4. Usually written 6.022 x 1023 and referred to as Avogadro’s
constant in honor of the Italian scientist Amedeo Avogadro.
5. Avagodro’s constant is defined as the number of particles, 6.022 x
1023 in exactly 1 mol of a pure substance.
6. The mole has been defined as the number of atoms in 12.00g of
carbon – 12.
a. Experiments have shown that 6.022 x 1023 is the number of
carbon-12 atoms in 12.00 g of carbon – 12.
b. One mole of carbon consists of 6.022 x 1023 carbon atoms with
an average mass of 12.01 amu.
Moles and grams are related
1. The mass in grams of 1 mol of a substance is called its molar mass.
a. E.g: 1 mol of C-12 atoms has a molar mass of 12.00 g.
b. A mole an element will usually include atoms of several isotopes.
c. The molar mass of an element in grams is the same as its average
atomic mass in a.m.u, listed in the periodic table.
2. The average atomic mass for carbon is 12.01 g.
Pictured: 1 mole of C or 12.00 g
Calculating with Moles
1. Since the amount of a substance and its mass are related, it is
useful to convert moles to grams, and vice versa.
2. Conversion factors are used to relate units.
3. Multiplying by a conversion factor is like multiplying by 1
because both parts of the conversion factor are always equal.
4. Example:
10 gumballs have a combined mass of 21.4
The relationship can be written as 2 equivalent conversion factors:
10 gumballs
21.4g
21.4g
10 gumballs
5. Now, we can use one of one these conversion factors to find the
mass of 50 gumballs, because mass increases in a predictable way
as more gumballs are added to the scale.
Steps for Calculating with Moles
1. What is the mass of exactly 50 gumballs?
Step 1: Given: 21.4g = 10 gumballs
Unknown: ?g = 50 gumballs
Step 2: Write down the conversion factor that converts number of
gumballs to mass.
The conversion factor you choose should have the unit you are solving
for (g) in the numerator and the unit you want to cancel (# of gumballs)
in the denominator:
21.4g
10 gumballs
Step 3: Multiply the number of gumballs by this conversion factor and
solve.
50 gumballs x 21.4g
= 107 g
10 gumballs
Practice - from page 132
1. What is the mass of exactly 150 gumballs?
150 gumballs x
21.4g
10 gumballs
= 321 g
2. If you want 50 eggs, how many dozens must you buy?
How many extra eggs do you have to take?
50 eggs x 1 dozen
12 eggs = 4.2 dozen
Therefore, 5 dozen eggs must be bought;
5 dozen x 12 eggs
1 dozen = 60 – 50 eggs = 10 extra eggs
3. If a football player is tackled 1.7 ft short of the end zone, how many
more yards does the team need to get a touchdown?
1.7 ft x 1 yd
3 ft = 0.57 yd
Relating amount to mass
1. Just as in the gumball example, there is also a relationship
between the amount of an element in moles and its mass in
grams.
2. (Refer to your book, page 132):
This relationship is graphed for iron nails because the amount of
iron and the mass of iron are directly related, the graph is in a
straight line.
3. An element’s molar mass can be used as if it were a conversion
factor. Depending on which conversion factor you use, you can
solve for either the amount of the element or its mass.
Converting moles (Amount) to grams (Mass)
E.g.: Determine the mass in grams in of 5.50 mol of Fe.
1. First you must know/find an element’s average atomic mass.
a. Given: amount of iron = 5.50 mol Fe
molar mass of Fe = 55.85 g/mol Fe
b. Unknown: mass of iron = ? g Fe
2. Write down the conversion factor that converts moles to grams.
a. The conversion factor you choose should have what you are trying
to find (grams of Fe) in the numerator and what you want to cancel
(moles of Fe) in the denominator.
b.
55.85 g Fe
1 mol Fe
3. Multiply the amount of iron by this conversion factor , and solve.
5.50 mol Fe x 55.85 g Fe
1 mole Fe = 307 g Fe
Practice: Converting Amount to Mass
Page 133:
(Work these on your paper):
What is the mass in grams of each of the following:
1. 2.50 mol of Sulfur, S
2. 1.80 mol of Calcium, Ca
3. 0.50 mol of Carbon, C
4. 3.20 mol of Copper, Cu
These will be worked out on the board if needed.
Converting Mass to Amount
1. Determine the amount of iron present in 352 g of Fe.
Step 1: Given: mass of iron = 352 g Fe
molar mass of iron = 55.85 g/mol Fe
Unknown: amount of iron = ? Mol Fe
Step 2 :
Write down the conversion factor that converts grams to moles.
The conversion factor you choose should have what you are
trying to find (moles of Fe) in the numerator and what you want
to cancel (grams of Fe) in the denominator.
1 mole Fe
55.85 g Fe
Step 3:
Multiply the mass of iron by this conversion factor, and solve.
352 g Fe x 1 mole Fe
55.85 g Fe = 6.30 mol Fe
Time for more practice!
Complete Section 4 Review p. 134 – all (1 – 8)
The next 40 elements:
1. Sc
2. Y
3. La
4. Ac
5. U
6. W
7. Hs
8. Ra
9. Ba
10. Ac
11. Rf
12. Fr
13. Os
14. Re
15. Ta
16. Ga
17. In
18. Tl
19. Pb
20. Bi
21. Po
22. Pd
23. Pt
24. Al
25. Hg
26. Bh
27. Hf
28. Ce
29. Th
30. Np
31. Pu
32. Nd
33. Pm
34. Sm
35. Eu
36. Am
37. Gd
38. Cm
39. Tb
40. Bk