Download Chapter 4 Atoms - LCMR School District

Chapter 4.2 Structure of Atoms
Parts of an Atom
All atoms have a centrally located Nucleus – center of atom, contains positive charged Protons and neutral
Neutrons(have no charge) and Negative charged Electrons that are orbiting the nucleus in an electron cloud.
Atoms have a neutral charge, the protons positive charge cancels out the electron’s negative charge. . The
opposite charges (Electric force) holds the atom together.
The arrangement of the electron determines the elements reactivity.
Ex. Unstable
stable
unstable
Atoms want to be stable. They either react and combine with other elements/molecules or become an ….
Ion - an atom with a positive or negative charge, more stable than neutral atoms. Many compounds are
formed from the attraction of these oppositely charged ions. Salt is one of them! NaCl
Atoms that have gained an electron are called anions Atoms that have lost an electron become a cation
(nonmetals elements). Atoms will fill almost full
(metals). Atoms with a few number of valence
valence energy levels with electrons and become
electrons lose their electrons and become cations.
anions. Because the atom has less protons than
Because the atom has more protons than electrons,
electrons, the “charge” on the atom is negative.
the “charge” on the atom is positive.
Key Terms to Know
Atomic number – each element has the same number of Protons; Elements have specific Atomic numbers,
Chlorine’s atomic number is 17
Average Atomic Mass – the average mass of an element, represents one mole of a that substance, Chlorine
has a average atomic mass of 35.453
Mass Number – the average atomic mass rounded to a whole number; represents the number of neutrons
and protons in the nucleus of an atom, Chlorine mass number is 35
Isotopes – Elements that have a different number of neutrons than the average.
The abundance of isotopes causes the average atomic mass to stray more from the mass number.
Molar Math
Unified Atomic Mass – unit of mass that represents 1/12 of the mass of a carbon 12 atom. (normal carbon)
Mole – SI base unit used to measure small particles; atoms, ions, or molecules
Avogadro’s Number – one mole of a substance equals 6.0223 x 1023 particles
Average atomic mass – one mole of substance equals the average atomic mass in grams, H = 1.00794 grams
Atom/Ion
Mass to Moles
Moles to Mass
Diatomic molecules
molecules
4.1 The Development of Atomic History
4th Century Greek Philosopher, Democritus, suggested Universe was made of invisible units called atoms –
“that which cannot be divided.” No evidence, so not many believed him.
Atomic Theory and Models
(grew as a series of models that developed from experimental evidence. As more evidence was collected,
the theory and models were revised)
1808- Dalton –
Billard Balls
1897- Thomson –
Cookie Dough
1911- Rutherford –
The Peach
John Dalton Atomic Theory
J.J. Thomson and Smaller Parts
Earnest Rutherford and the
Nucleus
*All elements are composed of
atoms that cannot be divided
*Atoms of same element are
exactly alike and have the same
mass, different elements have
atoms of different masses.
*Atoms of one element can’t be
changed into an atom of a
different element.
*Atoms cannot be created or
destroyed in any chemical
experiment, only rearranged.
*Compounds of atoms are
combined in specific ratios. Law
of Definite Proportions
1897- experimenting with
Cathode Ray Tubes, found atoms
contain negatively charged
particles, then reasoned due to
the neutral charge of atoms,
there must be a positive charge
particle to balance the charge.
1911- One of Thomson’s
students. Found evidence, “Gold
Foil experiment” to counter
Thomson’s model.
“negative charged particles
scattered throughout a ball of
positive charges”
negative charges particles later
became known as electrons
*Inferred that the positive
charged particles must be
clustered in a tiny region in its
center, called the nucleus.
*Atom was mostly empty space
with electron moving around the
nucleus in that space.
Modern Atomic Model
Bohr’s Model
A Cloud of Electrons
1913 - student of both Thomson
and Rutherford revised model
again.
1920’s – electrons do not orbit in
planet like orbits, instead they
can be found anywhere in a
cloudlike region around the
nucleus called the energy level.
The inner energy level can only
hold 2 electrons, and the larger
outer energy levels can hold, 8,
16, 32 respectively.
*electrons only have set energy
values, leading to orbit the
nucleus in specific orbital.
1932- James Chadwick discovered
the neutron and completed the
Modern Atomic Model.
*Neutron had no electrical
charge,
“Onion Model”
J.J. Thomson’s experiment
Rutherford’s – “Gold Foil” Experiment
Rutherfords Model – The Stadium Analogy
Why we colors!
Ground State vs. Excited State
Electrons jump from ground state to excited state when energy/photons are absorbed. Each particular
wavelength/frequency of light has a certain energy associated with it. This energy is directly related to the
“jump” to a new excited energy level. More energy absorbed, a more a jump. When the electron “jumps”
back to the ground state, the energy that was absorbed is now released. Our cones(color cells) in the eye
interpret the frequency and send a signal to our brain of for that color.
4.3 Electron Configuration
In the Bohr’s modern atomic model, electrons can only be found in certain energy levels, not between.
Electrons must gain energy to move to a higher level and lose energy moving to a lower level.
A new model, suggested that electrons act more like waves on a string, than like particles. Thus their exact
location cannot be predicted precisely. Scientists can predict the electrons orbital, the region where there is
the greatest probability finding an electron.
4 types of Orbital, each can hold a specific number of electrons.
Orbital
Axis orientation –each axis
holds 2 electrons
s
1
p
3
d
5
f
7
Total electrons
2
6
10
14
#refer to periodic table handout and ptable.com
Electron Transitions
We see because light is “reflected” back to our eyes, but how does this occur?
1. Photons are bundles of electromagnetic energy, each vibrating at a certain frequency. Higher the
frequency, higher the energy.
2. If this frequency is in visible spectrum, we can see it. The atom’s electrons absorbs this energy,
causing the electrons to get “excited” and jump from the “ground state” to a higher energy levels
within the atom, now called an “excited state.”
3. Then they immediately return to “ground state,” releasing the energy (photon) absorbed at a
similar frequency.
4. Our eyes are able to pick up this photons vibration, (frequency) and our brain interprets it as
color!! We see due the electrons jumping from a “ground state” to “excited state” and back. Every
wonder about it?