honors ch3 and ch4 - manasquan public schools...ch. 3.1 the atom is defined 400 b.c. the greek...
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HonorsCh3 and Ch4
Atomic History
and the Atom
Ch. 3.1The Atom is Defined
400 B.C. the Greek philosopher
Democritus said that the world was made
of two things:
– Empty space and tiny particles called
atoms
Ch3.1 What are atoms?
Atoms are the smallest
part of an element that
still has the element’s
properties.
By 1700’s chemistry was defined by 3 Laws:
Law of the Conservation of Mass
Law of the definite Proportions
Law of Multiple Proportions
Law of Conservation of Mass
Law of Definite Proportions/Composition
Substances contain atoms in the same
ratio of mass.
Law of Multiple Proportions
Early 1800’s John Dalton
Came up with first atomic
theory that is the basis for
today’s theory.
Early 1800’s John Dalton’s Theory-proven
1. Every element is made of tiny, unique particles called atoms– These atoms cannot be destroyed, but
instead rearrange during a chemical change
2. Atoms of different elements can join to form molecules in constant whole number ratios.
John Dalton’s Theory-disproved
Atoms cannot be broken down into smaller particles.
Atoms of the same element are exactly alike in mass
Ch3.21897 JJ Thomson
Used a cathoray tube to
examine if atoms were made
of charged particles.
Opposite charges attract
1897 JJ Thomson
Discovered atoms are
made up of particles
with negative charges,
but little mass.
Called them electrons.
Electrons
Positive charges, not
known as protons yet
Thomson Model
1911-Rutherford
Put together a team of
physicists to performed the
gold foil experiment
1 out of 8000 alpha
particles were repelled.
1913-Rutherford
The experiment led to the discovery that atoms are mostly empty space. (expected)
It also discovered that atoms contain a positive dense nucleus which contained most of the mass of the atom.
Rutherford Model
Inside the Nucleus Particles were discovered.
Positive protons (1836 times more massive
than electrons)
– These practices identified the type of atom.
Neutral neutrons (1837 times more massive
than electrons)
– Kept the protons from repelling by producing strong
nuclear forces
Ch 3.3 Parts of an atom
Nucleus
–Proton
–Neutron
Cloud
–Electron
Subatomic Particles
Charged Particles
Nucleus
center of an atom
positively charged
makes up 99.9% of the atom’s mass
Protons
Charge (+)
Mass is equal to
1 atomic mass unit (u)
–1/12 mass of Carbon atom
Neutrons
Charge (0 net) - neutral
Mass is equal to 1amu
Determine stability of
nucleus
Electrons
Charge is negative (-)
Mass is equal to 0.001 amu
Atomic Number
Identifies # of protons
Determines the type of atom because
no two elements can have same # of
protons.
Mass of a single atom
Mass Number
# of
protons
# of
neutrons
Mass
#
Isotopes
Any atoms having the same number of protons but different number of neutrons.
thus they have different mass numbers.
Springfield Isotopes
Isotopes
Average Atomic Mass
Atomic mass is the mass of all
isotopes of a particular element
averaged together
Calculating Average Mass is based
on the abundance of each isotope
Atoms vs Ions
All atoms have the same
number of protons and
electrons.
They are neutral.
Charges cancel each other out.
Atom vs Ions
Ions are charged particles.
Form when atoms lose or gain electrons.
Form in order to have a full outer shell
Two Types.
Cations
Positively charged ions.
Form when atoms lose
electrons.
Form from metal atoms
Cations
# of protons greater
than # of electrons
More (+) than (-)
Na AtomNa+ Cation
Anions
Negatively charged ions.
Form when atoms gain
electrons.
Form from nonmetal atoms
Anions
# of protons less
than # of electrons
More (-) than (+)
Cl atom
Cl- Anion
Ch4.1 The Duality of Light Led to a New View of the Atom
Light has characteristics of both waves and
particles
All forms of radiation travel at the same
maximum speed of 3.00 x108 m/s
Wave description
Different forms of light are defined by their
unique wavelengths (ƛ) and frequencies (v)
Speed of light (c)= ƛ v
Electromagnetic Spectrum
Electromagnetic spectrum
includes light at all possible
frequencies and wavelengths
↑ wavelength, ↓ frequency
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Gamma Rays
highest energy and frequency.
Nuclear radiation
X-Rays
high ionizing energy radiation
Used for imaging
Ultraviolet Light (UV)
Ionizing energetic radiation
Can cause skin cancer
when over exposed
Sun Protection (SPF)
7 or less→no protection
8→Extra protection but still permits tanning
15→Offers total protection from burning
30→Totally blocks UV
ROYGBIV
Visible spectrum
We see red and orange the best
Blue and violet are the hottest colors, and
emits the most energy, red the least and
coldest
White is all the colors combined, black the
absence
Infrared Light
Lower energy radiation
Night vision
Microwaves
Low energy radiation
Used to heat food
Also used in
telecommunication
Radio Waves
Have the Lowest
frequencies and highest
wavelengths
Includes FM, AM, and TVs
Radar
REALITY TV
The Particle Description
The Photoelectric effect:
– Emission of electrons when certain light hits
metal
Einstein theorized that light can be modeled as
a photons (particles of light)
– Each photon carries an unique quantum of energy
Maxwell Planck
Theorized that the energy given off by a photon is directly related to the frequency of the radiation emitted.
His equation Energy (E) = hv– h = planck’s constant = 6.626 x 10-24 J•s
Ch. 4.11920s Niels Bohr
Suggested that electrons move
around nuclei in set paths around
the nucleus. (solar system)
7. 1920s Niels Bohr
He said each path is a calculated
energy level
Atom’s electrons can jump to
different energy levels when
absorbing photons
Niels Bohr
States of Atoms
Ground State- An electron’s
lowest energy state or level
Excited State- An electron that is
energized will jump to a higher
energy state or level. Will last for a
short period.
– (resulting light production)
States of Atoms?
Ground State Excited state
Ground State Excited state back to ground
This is how light is produced
And Light production
Niels Bohr
Bohr noticed that different
atoms emitted different
radiation when excited.
Planck’s Flame Test
Spectroscopy
Emission (Line) Spectrum
Atoms and molecules are identified
by these spectrums.
What the Heck is Light?
Ch4.2 Quantum Theory
French Physicist De Broglie
Wave-Particle Duality of electrons
– He proposed that electrons can only exist at certain frequencies, hence the energy they release when excited
Today’s Theory
Werner Heisenberg Uncertainty Principle – It is impossible to determine an
electron’s exact position and speed at the same time.
Along with De Brogile, they disproved Bohr’s “definite orbit” assumption
Schrödinger’s Work
Electrons found in orbitals within different energy levels.
–a calculated region in an atom where there is a high probability of finding electrons.
This led to the Electron Cloud model and Quantum Numbers
Modern Atomic Cloud Model
History Review
Quantum #’s
Principal (n) = number of energy level (1-7)
Angular momentum (l) = sublevels: 0,1,2,or 3
Magnetic (m) = the orbital # (-,0,+)
Spin (s) = electron spin (1/2, -1/2)
Pauli Exclusion Principle
• No two atoms will have the same
configuration or set of quantum #’s
Energy levels
1st level holds 2 e- (s)
2nd level holds 8 e- (s,p)
3rd level holds 8 or 18e- (s,p,d)
4th level holds 18 or 32e- (s, p d, f)
Outer (valence) level holds up to 8 e-
(s, p)
Making Bohr Models
Bohr’s Model
Blue dots represent electrons
Rings represent energy Level,NOT orbit (path)
ELECTRON
CONFIGURATIONS
Ch4.2
Electron Energy Levels, SUBLEVELS, and Orbital's
Electron Configuration
The rows (periods) of the periodic table tell you the energy level
There are 4 sublevels – s,p,d,f
Electron Configuration
Each sublevel contains a different number of orbitals.
S =1, p = 3, d= 5, and f = 7
Orbital is represented by:
Each orbital holds only two electrons with opposite spin.
Writing Atomic Electron Configurations
Three ways of writing configurations.
1. orbital box notation.
2. Electron Configuration (spdf) Notation
3. Noble Gas Configuration
Orbital Notation Rules:
1. Aufbau Principle- electrons are added one
at a time starting at lowest energy level
2. Hund’s Rule - When filling orbitals with in
the same sublevel, fill one electron into each
box before pairing electrons.
Orbital Notation
Arrowsdepictelectronspin
ORBITAL BOX NOTATIONfor He, atomic number = 2
1s
21 s
3. Hund’s Rule1.
2P
S,p,d,f Notation
11 s
value of nsublevel
no. of
electrons
for H, atomic number = 1
Phosphorus
Group 5A
Atomic number = 15
1s2 2s2 2p6 3s2 3p3
[Ne] 3s2 3p3
1s
2s
3s
3p
2p
Aluminum
Group 3A
Atomic # = 13
1s2 2s2 2p6 3s2 3p1
[Ne] 3s2 3p1
All Group 3A elements have [core] ns2 np1
configurations where n is the period number.
1s
2s
3s
3p
2p