Chapter 2

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Transcript Chapter 2

Chapter 2
Atoms and Elements
Law of Conservation of Mass


in a chemical reaction, matter is neither created nor destroyed
total mass of the materials you have before the reaction must equal the
total mass of the materials you have at the end
◦ total mass of reactants = total mass of products
7.7 g Na
+
11.9 g Cl2

19.6 g NaCl
Conservation of Mass and the Law of
Definite Proportions
Law of Definite Proportions (John Dalton 1766-1844)
-Different samples of a pure chemical substance always contain the
same proportion of elements by mass.
-Compound always forms from a fixed ratio of its element
How many Hydrogen atoms are there?
How many Oxygen atoms are there?
What is the ratio between Hydrogen atoms to
Oxygen atom?
Now using the mass of Hydrogen and Oxygen to show these results are
consistent with the law of definite proportion
Proportions in Sodium Chloride
a 100.0 g sample of sodium
chloride contains 39.3 g of
sodium and 60.7 g of chlorine
mass of Cl 60.7 g

 1.54
mass of Na 39.3 g
a 200.0 g sample of sodium
chloride contains 78.6 g of
sodium and 121.4 g of chlorine
mass of Cl
 ???
mass of Na
a 58.44 g sample of sodium
chloride contains 22.99 g of
sodium and 35.44 g of chlorine
mass of Cl
 ???
mass of Na
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The Law of Multiple Proportions and
Dalton’s Atomic Theory
Law of Multiple Proportions: If two elements (A and B) form more
than one compound (AB, AB2, AB3 …) then the ratios of the masses of
the second element which combine with a fixed mass of the first
element will be the ratio of small whole number
- Show how chemical formulas are put together
nitrogen monoxide (NO): 7 grams nitrogen per 8 grams oxygen
nitrogen dioxide (NO2): 7 grams nitrogen per 16 grams oxygen
Insert Figure 2.2 p37
Example


Lead forms two compounds with oxygen as shown:
PbO:
2.98 g Pb, 0.461 g O
PbO2:
9.89 g Pb, 0.763 g O
Show these results consistent with the law of definite
proportions
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Dalton’s Atomic Theory
Dalton proposed a theory of matter based on it having
ultimate, indivisible particles to explain these laws
1)
Each element is composed of tiny, indestructible
particles called atoms
2)
All atoms of a given element has the same mass and
other properties that distinguish them from atoms of
other elements
3)
Atoms combine in simple, whole-number ratios to form
molecules of compounds
4)
In a chemical reaction, atoms of one element cannot
change into atoms of another element
◦ they simply rearrange the way they are attached

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J.J. Thomson (1897)
believed that the cathode ray was composed of tiny particles with
an electrical charge
 designed an experiment to demonstrate that there were particles by
measuring the amount of force it takes to deflect their path a given
amount
◦ like measuring the amount of force it takes to make a car turn

anode
cathode
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Thomson’s Experiment
investigate the effect of placing an electric field around
tube
(1) charged matter is attracted to an electric field
(2) light’s path is not deflected by an electric field
+++++++++++
cathode
anode
(+)
(-)
-------------
9
Power Supply
+
Thomson’s Results
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the cathode rays are made of tiny particles
these particles have a negative charge
◦ because the beam always deflected toward the + plate
Particles that make cup cathode rays are 2,000 times smaller than a
Hydrogen
the amount of deflection was related to two factors, the charge
and mass of the particles
every material tested contained these same particles
the only way for this to be true is if these particles were pieces of
atoms
◦ apparently, the atom is not unbreakable
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Millikan’s Oil Drop Experiment
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Millikan's experiment involved
measuring the force on oil
droplets in a glass chamber
sandwiched between two
electrodes, one above and one
below. With the electrical field
calculated, he could measure the
droplet's charge,
the charge/mass of these
particles was -1.76 x 108 C/g
◦ the charge/mass of the
hydrogen ion is +9.58 x 104
C/g
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Thomson’s Plum Pudding Atom
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The Plum Pudding Atom
◦ the mass of the atom is due to the mass of the electrons within it
◦ electrons are the only particles in Plum Pudding atoms
◦ the atom is mostly empty space
◦ cannot have a bunch of negatively charged particles near each other as
they would repel
the structure of the atom contains many negatively charged electrons
these electrons are held in the atom by their attraction for a positively
charged electric field within the atom
◦ there had to be a source of positive charge because the atom is neutral
◦ Thomson assumed there were no positively charged pieces because
none showed up in the cathode ray experiment
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Rutherford’s Results
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Over 98% of the a particles went straight through
About 2% of the a particles went through but were deflected by large
angles
About 0.01% of the a particles bounced off the gold foil
◦ “...as if you fired a 15” cannon shell at a piece of tissue paper and it
came back and hit you.”
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Plum Pudding
Atom
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•
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a few of the
a particles
do not go through
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•
•
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•
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if atom was like
a plum pudding,
all the a particles
should go
straight through
•
Nuclear Atom
.
.
.
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most a particles
go straight through
some a particles
go through, but are deflected
Rutherford’s Interpretation –
the Nuclear Model
1)
2)
3)
4)
The atom contains a tiny dense center called the nucleus
 the amount of space taken by the nucleus is only about
1/10 trillionth the volume of the atom
The nucleus has essentially the entire mass of the atom
 the electrons weigh so little they give practically no mass to
the atom
The nucleus is positively charged
 the amount of positive charge balances the negative charge
of the electrons
The electrons are dispersed in the empty space of the atom
surrounding the nucleus
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Structure of the Atom
Rutherford proposed that the nucleus had a particle that
had the same amount of charge as an electron but
opposite sign
◦ based on measurements of the nuclear charge of the
elements
 these particles are called protons
◦ charge = +1.60 x 1019 C
◦ mass = 1.67262 x 10-24 g
 since protons and electrons have the same amount of
charge, for the atom to be neutral there must be equal

numbers of protons and electrons
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Relative Mass and Charge

we generally talk about the size of charge on atoms by
comparing it to the amount of charge on an electron, which
we call -1 charge units
◦ proton has a charge of +1cu
◦ protons and electrons have equal amounts of charge, but opposite
signs

we generally talk about the mass of atoms by comparing it to
1/12th the mass of a carbon atom with 6 protons and 6
neutrons, which we call 1 atomic mass unit
◦ protons have a mass of 1amu
◦ electrons have a mass of 0.00055 amu, which is generally too small to
be relevant
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Some Problems

How could beryllium have 4 protons stuck together in the
nucleus?
◦ shouldn’t they repel each other?

If a beryllium atom has 4 protons, then it should weigh 4 amu;
but it actually weighs 9.01 amu! Where is the extra mass coming
from?
◦ each proton weighs 1 amu
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There Must Be Something Else
There!
to answer these questions, Rutherford proposed that there was
another particle in the nucleus – it is called a neutron
 neutrons have no charge and a mass of 1 amu
◦ mass = 1.67493 x 10-24 g
 slightly heavier than a proton
◦ no charge

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Summary: The Structure of the Atom
Atom: The basic unit of an element that can enter into chemical combination
They are made of even smaller particles called subatomic particles.
Electrons
Protons
Neutrons
Subatomic Mass
Particle
Proton
g
Mass Location Charge Symbol
amu
in atom
1.67262 1.00727 nucleus
+1
p, p+, H+
-1
e, e-
0
n, n0
x 10-24
Electron 0.00091 0.00055
x 10-24
empty
space
Neutron 1.67493 1.00866 nucleus
x 10-24
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Elements


each element has a unique number of protons in its nucleus
the number of protons in the nucleus of an atom is called the atomic
number
◦ the elements are arranged on the Periodic Table in order of their
atomic numbers
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Sources of Some Element Names
Some elements are named for planets,
mythological figures, minerals, colors,
scientists, and places.
A symbol
•represents the name of an element.
•consists of 1 or 2 letters.
•starts with a capital letter
•With few elements use their own
special symbols
1-Letter Symbols
C
carbon
N
nitrogen
S
sulfur
Diatomic Molecule: A molecule
that consists of two atoms (N2,
O2, H2, F2, Cl2, Br2, I2).
The diatomic molecules
(H2,N2,O2,F2,Cl2) are gaseous in
their elemental forms at normal
temperatures. Br2 and Hg are
liquids in their elemental forms
at normal temperatu
2-Letter Symbols
Co cobalt
Ca calcium
Au
gold
Mg magnesium
2
2
Mendeleev and the Periodic Table
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order elements by atomic mass
saw a repeating pattern of properties
Periodic Law – When the elements are arranged in order of
increasing atomic mass, certain sets of properties recur
periodically
put elements with similar properties in the same column
used pattern to predict properties of undiscovered elements
where atomic mass order did not fit other properties, he reordered by other properties
◦ Te & I
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Groups: The elements in a vertical column of the periodic table.
Period: The elements in each horizontal row of the periodic table.
Alkali Metals: The Group 1A elements (Li,Na,K,Rb,Cs,Fr).
Alkaline Earth Metals: The Group 2A elements (Be,Mg,Ca,Sr,Ba,Ra).
Halogens: The nonmetallic elements in Group 7A (F,Cl,Br,I,At).
Noble Gases: The nonmetallic elements in Group 8A
(He,Ne,Ar,Kr,Xe,Rn).
Transition Metals: Group of elements in the middle of the periodic
table.
Metals
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solids at room temperature, except Hg
reflective surface
◦ shiny
conduct heat
conduct electricity
malleable
◦ can be shaped
ductile
◦ drawn or pulled into wires
lose electrons and form cations in reactions
about 75% of the elements are metals
lower left on the table
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Sulfur, S(s)
Nonmetals
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found in all 3 states
poor conductors of heat
poor conductors of electricity
solids are brittle
gain electrons in reactions to become anions
upper right on the table
◦ except H
Bromine, Br2(l)
Chlorine, Cl2(l)
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Metalloids
show some properties of
metals and some of
nonmetals
 also known as
semiconductors

Properties of Silicon
shiny
conducts electricity
does not conduct heat well
brittle
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Structure of the Nucleus


Soddy discovered that the same element could have atoms with different
masses, which he called isotopes
The observed mass is a weighted average of the weights of all the
naturally occurring atoms
◦ the percentage of an element that is 1 isotope is called the isotope’s
natural abundance
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Isotopes

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all isotopes of an element are
chemically identical
◦ undergo the exact same
chemical reactions
all isotopes of an element have
the same number of protons
isotopes of an element have
different masses
isotopes of an element have
different numbers of neutrons
isotopes are identified by their
mass numbers
◦ protons + neutrons
Atomic Number
 Number of protons
Z
• Mass Number
 Protons + Neutrons
 Whole number
A
• Abundance = relative
amount found in a sample
•
Isotopic symbols
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Examples
1. How many protons, electrons and neutrons are
present in an atom of
2. Write isotopic symbols in both forms for Selenium
isotope with 40 neutrons
3. An atom has 32 electrons and 38 neutrons. What is
its mass number and what is the element?
Isotopes and atomic mass
20
10
90.48%
Ne
21 Ne
10
10
11
21
0.27%
22 Ne
10
10
12
22
9.25%
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Atomic Mass

we previously learned that not all atoms of an
element have the same mass
◦ isotopes
we generally use the average mass of all an
element’s atoms found in a sample in calculations
 we call the average mass the atomic mass

Atomic Mass   fractional abundance of isotope n  mass of isotope n
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Mass Spectrometry
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masses and abundances of isotopes are
measured with a mass spectrometer
atoms or molecules are ionized, then
accelerated down a tube
their path is bent by a magnetic field,
separating them by mass
◦ similar to Thomson’s Cathode Ray
Experiment
a mass spectrum is a graph that gives the
relative mass and relative abundance of each
particle
relative mass of the particle is plotted in the
x-axis
relative abundance of the particle is plotted
in the y-axis
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Example

Magnesium has three naturally occurring isotopes with masses of 23.99
amu, 24.99 amu, and 25.98 amu and natural abundances of 78.99%,
10.00% and 11.01% respectively. Calculate the atomic mass of
magnesium

Lithium has two naturally occurring isotopes: lithium-6 and lithium-7. If
the average atomic mass of lithium is 6.941 amu, which isotope is the
most abundant? How do you know?
Reacting Atoms
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when elements undergo chemical reactions, the reacting elements do not
turn into other elements
◦ Dalton’s Atomic Theory
since the number of protons determines the kind of element, the number
of protons in the atom does not change in a chemical reaction
however, many reactions involve transferring electrons from one atom to
another
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Charged Atoms
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when atoms gain or lose electrons, they acquire a charge
charged particles are called ions
when atoms gain electrons, they become negatively charged ions, called
anions (Cl-)
when atoms lose electrons, they become positively charged ions, called
cations (Na+)
ions behave much differently than the neutral atom
◦ e.g., The metal sodium, made of neutral Na atoms, is highly reactive
and quite unstable. However, the sodium cations, Na+, found in table
salt are very nonreactive and stable
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Ionic Charge from Group Numbers
•
The charge of a positive ion is equal to its Group number.
Group 1A(1) = 1+
Group 2A(2) = 2+
Group 3A(3) = 3+
•
The charge of a negative ion is obtained by subtracting 8 or 18
from its Group number.
Group 6A(16) =
6-8
or 16 - 18
= 2-
= 2-
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Atomic Structures of Ions
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Nonmetals form anions
For each negative charge, the ion has 1 more electron than the neutral
atom
◦ F = 9 p+ and 9 e-, F ─ = 9 p+ and 10 e-
•Metals form cations
•For each positive charge, the ion has 1 less electron than the
neutral atom
Na atom = 11 p+ and 11 e-, Na+ ion = 11 p+ and 10 e-
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Examples

Predict the charge and identify the number of proton and electron from
that ion
◦ Mg 
◦ Al 
◦ O
◦ P
Collection Terms
A collection term states
a specific number of items.
• 1 dozen donuts
= 12 donuts
• 1 ream of paper
= 500 sheets
• 1 case
= 24 cans
In chemistry, how do chemists know the number of atoms
in an element or compound?
2
A Mole of Atoms
A mole is
• a unit of measurement used in chemistry to express amounts
of a chemical substance, the same number of particles as there
are carbon atoms in 12.0 g of carbon.
• a collection term “dozen”
1 mole = NA = 6.022 x 1023 of anything
Avogadro’s Number = 6.0221421 x 1023
3
Relationship Between Moles and Mass


The mass of one mole of
atoms is called the molar
mass
The molar mass of an
element, in grams, is
numerically equal to the
element’s atomic mass, in
amu
E.g
1 H atom = 1.01 amu
1 mol H = 1.01g
6.022 x 1023 atoms of H = ???? g
42
Examples
Give the molar mass for each
A. 1 mole of Li atoms =
________ g
B. 1 mole of Co atoms =
________g
C. 1 mole of S atoms =
________g
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Converting between mass, moles and
atoms
gC
mol C
mol C
gC
gC
mol C
atoms
Examples

Assuming all pennies are pure copper and each has a mass of 2.5 g
◦ Without doing calculation, determine the number atoms of copper
present in 1 mole
◦ How many pennies does it take to make a mole?
Examples

Calculate the moles of carbon in 0.0265 g of pencil lead

Calculate the mass (in grams) of 0.473 moles of titanium

In a 3.0 moles of O2 molecules
◦ How many oxygen molecules are there ?
◦ How many oxygen atoms are there?