Download UNIT 5 - ATOMIC THEORY: THE NUCLEAR MODEL OF THE ATOM

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UNIT 5 - ATOMIC THEORY: THE NUCLEAR MODEL OF THE ATOM
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Dalton’s Atomic Theory
1) Each element is made up of tiny, individual particles called atoms.
2) Atoms are indivisible; they cannot be created or destroyed.
3) All atoms of each element are identical in every respect.
Dalton’s Atomic Theory
4) Atoms of one element are different from atoms of any other element.
5) Atoms of one element may combine with atoms of another element,
usually in the ratio of small whole numbers, to form chemical compounds.
Dalton’s theory led to the Law of Multiple Proportions.
The Law of Multiple Proportions
Elements may combine in more than one proportion to form more than one
compound.
Examples...
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Subatomic Particles
Dalton’s theory about the atom being indivisible was challenged with the
discovery of subatomic particles within an atom.
Three types we’ll look at:
1) Electrons
Electrons are negatively - charged particles that were discovered in 1820 by
Michael Faraday.
properties: it has a charge of -1 and a mass of 9.107 x 10-28 grams.
Electrons
Thompson proposed his raisin bun model of the atom - Electrons are evenly
distributed inside the spherical positive part of the atom.
2) Protons
Protons are positively - charged particles that were discovered in 1919 by
Ernest Rutherford.
properties: it has a charge of +1, and has a mass that is 1836 times greater
than the mass of an electron (1.67 x 10-24 g).
3) Neutrons
Neutrons are electrically neutral particles discovered in 1932 by James
Chadwick.
A neutron’s mass is approximately equal to the mass of a proton (1.67 x
10-24 grams).
A summary of subatomic particles is given in Table 5.1 on page 117.
Ernest Rutherford (1871 - 1937)
Rutherford formed the nuclear model of the atom.
Overhead: Figures 5.3 & 5.4, p. 118
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An atom consists of a very small, positively charged nucleus, in which most
of the mass of the atom is concentrated, surrounded by the electrons
necessary to produce an electrically neutral atom.
The Nuclear Model of the Atom
At the time, Rutherford only knew about electrons.
Electrons circled the positively charged nucleus like the planets orbit the sun
- the planetary model of the atom.
The Nuclear Model of the Atom
Problems with Rutherford’s model:
A) If the electron is stationary, the atom should disappear because the
electron would be electrostatically drawn in to the nucleus.
B) If the electron orbits the nucleus, it will have to change direction a lot. A
change in direction means acceleration, which involves the use of energy.
If this is so, the electron will again spiral into the nucleus.
Neils Bohr (1885 - 1962)
Bohr was a Danish physicist ($16,000 Millionaire question!!)
Bohr model
All of the atom’s positive charge and nearly all of its mass are contained in
an extremely dense nucleus
Negatively charged electrons of very small mass travel in orbits around the
nucleus.
The Bohr model
Bohr proposed that the energy of an electron in an atom cannot vary
continuously, but is instead quantized - at any instant, the electron may
have one of several possible energies, but at no time may it have an
energy between them.
Where can the electrons be?
The electron can orbit the nucleus at certain specified distances, but it is
never found between them.
example: a ladder, or steps
Normally, electrons are in the Ground State - the condition when all
electrons in an atom occupy the lowest possible energy levels.
Where can the electrons be?
An electron can move between orbits by quantum jumping or quantum
leaping: the electron disappears from one orbit and reappears in another.
Excited State - the condition at which at least one electron in an atom is at
an energy level above ground state.
Where can the electrons be?
An electron in an excited state is unstable - will fall back to the ground state
and release energy.
This energy release shows up as a line in the spectrum of the element.
Overhead - figure 10.3
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Why was Bohr’s model so good?
1) He suggested a reasonable explanation for the atomic line spectra in
terms of electron energies.
2) He introduced the idea of quantized electron energy levels in the atom
Problems with the Bohr model
1) It only fit the hydrogen atom.
2) According to the Law of Conservation of Energy, an electron circling a
nucleus should lose energy and crash into the nucleus.
For the next 13 years, people worked to improve these faulty parts, leading
to the Quantum Mechanical Model of the Atom.
Isotopes
It was discovered that all atoms of an element are not identical - some
atoms have more mass than other atoms of the same element.
Atomic number (Z) - every atom of an element has the same number of
protons.
Isotopes
Atoms are electrically neutral, therefore the same number of electrons must
be present.
What’s left to cause the difference in masses??
Answer: Neutrons
Isotopes
Isotopes - atoms of the same element that have different masses
(different numbers of neutrons).
Mass number (A) = # protons (Z) + # neutrons
The name of an isotope is its elemental name followed by its mass number.
Examples: carbon, oxygen, hydrogen
Atomic Mass
We recognize that a sample of a pure element contains atoms that have
different masses
The masses of atoms are expressed in atomic mass units (amu)
“Atomic Mass Scale” handout...
The Periodic Table
Dmitri Mendeleev and Lothar Meyer were cofounders of the periodic table in
1869.
They found that certain properties repeat at regular intervals when the
elements are arranged according to their atomic masses.
The Periodic Table
However, in order for all elements to fall into the proper groups, it was
necessary to switch a few of them, which interrupted the orderly increase
in atomic masses.
Reasons: there were errors in atomic weights as they were known in 1869
and, more importantly, it was later found that the correct ordering
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property is the atomic number (Z), not the atomic mass.
The Periodic Table
Horizontal rows on the table are periods.
Vertical columns on the table are groups, or chemical families.
Figure 5.6 - Period and group numbering, etc.
The Periodic Table
Main group elements - A groups (1,2 and 13-18)
Transition elements (transition metals) - B groups (3-12)
Metals and nonmetals...
Elemental Symbols and the Periodic Table
Pg. 128: Figure 5.8a - gives the name, symbol, and atomic number of the
elements whose names and symbols are to be learned.
Figure 5.8b - shows where they are on the periodic table.
Valence Electrons
Are the outermost and most reactive electrons in an element.
Elements in the same column, or group, of the periodic table have the same
number of valence electrons, and similar chemical properties.
Valence Electrons
Examples: Lithium, sodium, potassium
Their valence shell contains only 1 electron (Group 1)
All react similarly with water
Valence Electrons
Atoms that have a full set of eight valence electrons are said to have a full
octet of electrons.
These atoms tend to be mostly unreactive.
Example: the noble gases [Group 18 (8A)]
Elements strive to attain a full octet of electrons and the same electron
configuration as (become isoelectronic with) the nearest noble gas.
Lewis Symbols (Electron dot diagrams)
The symbol of the element is surrounded by that number of dots that
matches the number of valence electrons.
Examples… Na, C, N, Cl
Note: Don’t pair up electrons until there is one electron on each side of the
symbol.
Trends in the Periodic Table
ION - an atom or group of atoms that has an electrical charge because of a
difference in the number of protons and electrons.
eg. Na+, ClIONIZATION ENERGY - the energy required to remove one electron from a
neutral gaseous atom.
Ionization Energy
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Figure 10.12
Ionization energies increase as atomic number increases within a period.
Why?
Larger +ve nucleus, electrons same distance away.
Ionization energies decrease as atomic number increases within the group.
Why?
Electrons are further away.
Ionization Energy
The amount of energy required to remove a third electron is higher than
that required to remove a second electron, which is higher than the
amount of energy required to remove the first electron
Chemical Families
Elements in the same group on the periodic table have similar chemical
properties.
Four chemical families that we will look at:
1) the alkali metals - Group 1A (1)
2) the alkaline earths - Group 2A (2)
3) the halogens - Group 7A (17)
4) the noble gases - Group 8A (18)
Figure 10.13
1) The alkali metals
Group 1, therefore number of valence electrons is 1. Charge on the ions?
1+.
The reactivity of the element increases as you go down the group. Why?
Trends in properties of alkali metals:
A) Density increases as atomic number increases.
B) Boiling and melting points generally decrease as you go down the
periodic table.
2) The alkaline earths
Group 2, therefore number of valence electrons is 2. Charge on the ions?
2+.
Reactivity increases as you go down the group.
3) The halogens
Group 17, therefore number of valence electrons is 7 (17-10). Charge on
the ions?
1-.
Reactivity decreases as you go down a group. Why?
The larger the atom, the less easy it is to hold onto the new electron.
Density, melting point, and boiling point all increase steadily with increasing
atomic number.
4) The noble gases
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Group 18, therefore number of valence electrons is 8 (18-10). (Except for
He - only has 2 valence electrons) Ion charges?
none
Are reluctant to react, because they already have a full octet of electrons,
and don’t want any more. Helium is fine with 2 valence electrons.
Density, melting point, and boiling point all increase as you go down the
group.
Hydrogen
Valence electrons: 1.
Is neither an alkali metal nor a halogen, although it shares properties with
both.
Can lose an electron to form H+, or gain an electron to form H-.
Atomic Size
Size of atoms increases from right to left, and from top to bottom.
Figure 10.14
Why? (2 reasons)
1) Distance of valence electrons from the nucleus.
2) Nuclear charge.
Metals and Nonmetals
Figure 10.15 - staircase boundary
Generally, an element is a metal if it can lose one or more electrons to
attain a noble gas - like electron configuration
Ions of metals have positive charges.
Metallic character increases as you go down a group, but decreases as you
go across a period from left to right. Why?
Atom gets bigger, electron is further away.
Metals and Nonmetals
As you go down a group the atom gets bigger - it’s easier to lose one
electron as the atom gets bigger and the electron is further away.
As the period increases so does the amount of +ve charge in the nucleus tighter hold on the valence electrons.
Questions: Units 5 and 10
Unit 5 - You should be able to do the target check questions and examples
throughout the chapter, as well as the end - of - chapter questions that
begin on pg. 130. The blue - numbered questions have answers provided
at the end of the chapter.
Unit 10 - questions begin on page 288
Specific questions: #3, 4, 5, 6, 8, 28, 31, 32, 33, 35, 36, 38, 39, 40, 41,
43, 44, 45,
47 (a-e, k-p), 50