Download Atomic Structure Review–Honors

Matter, and more!
Matter
• Anything that has MASS and takes up SPACE
How is Matter classified?
• 1) Pure Substances
• 2) Mixtures
1) Pure Substances
• Composition remains the same, does
not depend on a sample = Fixed
composition
• Homogenous—same throughout.
• Example: Compound (NaCl) or Element
(Fe)
2) Mixtures
• 2+ substances combined together
• Substances do not change their properties or
name.
• Able to be separated, not chemically
combined.
• Possess a combination of properties based
on the substances present.
Types of Mixtures
1) Homogenous
– Uniform composition
– Also known as “true solutions”
– Ex. Salt-water
2) Heterogeneous
– No uniform composition
– Can easily see the different components of
the mixture
– Ex. Italian dressing
Properties of Matter
1) Chemical
– Ability to go through changes resulting in a
different substance
– The substance is no longer the same, different
identity
– Ex. Burning
2) Physical
– Observed or measured property
– Substance identity is not changed
– Ex. melting point, boiling point, density
Practice
Extensive vs. Intensive Properties
• Extensive Property
– Dependent on the quantity of matter
– Ex. Mass
• Intensive Property
– Does NOT depend on matter quantity
– Ex. Density
The Atom
Law of Conservation of
Mass/Matter
• Matter cannot be created or
destroyed
• Total mass is constant in chemical
reactions.
• Originated with Antoine Lavoister
(1700s)
– Quantitative mass data of
reactants and products in
mercury oxide decomposition.
Terminology
• Element– basic unit of a substance, contain
only ONE type of atom, represented by
symbol.
Example: Ag, only contains Ag atoms.
• Atom—smallest particle of an element that
still contains element properties.
– Example: One atom of Au, cannot have a
smaller particle of gold and still be gold.
Compound vs. Molecule
• Compounds:
– more than one element
– elements combined in definite proportions
• Molecule:
– Smallest unit of a compound that still retains the
properties of the compound.
Dalton Atomic Theory (cont.)
• All matter made of
atoms
• Atoms of an element
have the same size,
mass, etc.
• Different atoms have
various sizes, mass,
etc.
• Atoms cannot be
divided, destroyed, or
created.
• Atoms rearrange in
chemical reactions.
John Thomson
• 1897
• Cathode-Ray experiments.
• Discovered the electron
particle and its possible
charge (-).
• Determined ratio
between mass and charge
of an electron
Robert Millikan
• Oil drop
experiments.
• 1909, American
• Found the mass of an
electron (VERY small) with
Thompson’s data
• Currently, mass of electron =
9.109 x 10-31kg
• Discovered electron charge
– e = -1.602 x 10-19 C
Early Models of the Atom
Thompson
• Must be a balance between negative and positive
charges
• “Raisin-Pudding” model
• Uniform distribution of positive charge
– Positive cloud with stationary electrons
Early Models of the Atom
Rutherford
•
•
•
•
Mostly empty space
Small, positive nucleus
Contained protons
Negative electrons
scattered around the
outside
Early Models of the Atom
Bohr
• 1913—hydrogen atom
structure
• Physics + quantum theory
• Electrons move in definite
orbits around the positively
charged nucleus—planetary
model
• Does not apply as atoms
increase in electron number
Heisenberg’s Uncertainty
Principle
• Electron’s location and direction cannot
be known simultaneously
• Electron as cloud of negative charge
Modern Model of the Atom
The electron cloud
• Sometimes called the
wave model
• Electron as cloud of
negative charge
• Spherical cloud of varying
density
• Varying density shows
where an electron is
more or less likely to be
How did we discover electron
arrangement in an atom?
ELECTROMAGNETIC
RADIATION ! ! !
Electromagnetic Waves
• Produced from electric charge movement
• Changes within electric and magnetic fields
carried over a distance
• No medium needed
Types of Spectra
• What is a spectra?
– Spectrum– white light/radiation split into different
wavelengths and frequencies by a prism
• Continuous spectrum
– No breaks in spectrum
– Colors together
• Line spectrum
– Line pattern emitted by light from excited atoms of
a particular element
– Aided in determining atomic structure
Line Spectrum
• Pattern emitted by light from excited atoms
of an element
• Specific for each element
• Used for element identification
THEREFORE ………
• Light is in the form of electromagnetic waves
• Photons can resemble particles
• Gave raise to the possibility of thinking about
wave AND particle qualities of subatomic
particles (electron)
Atomic Structure
• Nucleus
– Protons
– Neutrons
• Electrons
Describing Atoms
• Atomic Number = number of protons
• In a neutral atom, the # of protons = the
# of electrons
• Atomic Mass= the number of protons +
the number of neutrons
Isotopes
• The number of protons for a given atom
never changes.
• The number of neutrons can change.
• Two atoms with different numbers of
neutrons are called isotopes
• Isotopes have the same atomic #
• Isotopes have different atomic Mass #’s
Ions
• An atom that carries an electrical charge is
called an ion
• If the atom loses electrons, the atom
becomes positively charged.
• If the atom gains electrons, the atom
becomes negatively charged
PEN Method for--•
•
•
•
F
Li +1
Cl -1
Be
Light
• Composed of small energy packets (photons)
• Quantum = minimum amount of energy lost/gained
by atom
• Atoms can absorb or give off this energy
Energy States in an Atom
• Atoms can gain or loss energy.
• Specific energy states within an atom.
– Can be counted
– Ground State = lowest energy state
– Excited State = higher energy level than ground,
gained energy
So, where does the Bohr Model fit
in?
• Electrons orbit around the nucleus at different energy
levels/orbits.
• Electron’s energy level = orbit level where electron is
located.
• Light absorption = electron moves from a state of low
energy to high energy. “becomes excited”
• Light Emitted = electron falls from an “excited” state
of energy to a lower energy level.
Main Energy Levels/Electron Shell
• n=1
– Holds 2 electrons
• n=2
– Holds 8 electrons
• n=3
– Holds 18 electrons
• n=4
– Holds 32 electrons
Energy sublevels
• Within the main energy level.
•
•
•
•
S = 1 orbital, can hold 2 electrons
p = 3 orbitals, can hold 6 electrons
d = 5 orbitals, can hold 10 electrons
f = 7orbitals, can hold 14 eletrons
Quantum Theory
• Treats electron’s location as wave property
• Defined by quantum numbers
• Orbitals have different energies
• Quantum numbers
– Provide information about size, shape, and orientation
of atomic orbitals
– Define atomic orbitals from general to specific
Quantum Mechanical Model
• Opposite charges attract, electrons are
attracted to the nucleus of an atom
– Takes a LOT of energy to keep electrons away from
the nucleus.
• Electrons are found at differing lengths from
the nucleus and can only be present in certain
locations
Principal Quantum Number (n)
• Determines orbital size and electron energy
• Same as “n” value/orbital in Bohr model
• Positive whole number, NOT 0
• Shells – orbitals with same value
• n = 1, 2, 3, 4, etc.
Orbital Angular Momentum
Quantum Number (l)
• Defines orbital shape for a particular region of atom
• Think of as “subshell”
• Energy sublevels—within the main energy level
–
–
–
–
–
s = 1 orbital, can hold 2 electrons
p = 3 orbitals, can hold 6 electrons
d = 5 orbitals, can hold 10 electrons
f = 7 orbitals, can hold 14 electrons
Orbital Shapes
• s orbital
– 1 possible orbital orientation, spherical shape
– n value determines size
– Charge cloud found near center, likely electron
location
• p orbital
– 3 possible orbital orientations, dumbbell shape
– pX, py, pz
l
Orbital/Subshell
0
s
1
p
2
d
3
f
How do you specify orbitals?
• 2p
• 4f
Electron Configuration
• Shorthand method for representing electrons’
distribution in orbitals within subshells
• All orbitals have the same energy level—digenerate
• Orbitals– mathematical expressions of probability of
electron’s location
• Electrons occupy orbitals in a way that gives LOWEST
energy state
Orbital Diagrams
• Visual representation of electron
configuration
• Represents electrons’ spins (, )
Aufbau Principle
• Electrons occupy the LOWEST energy orbital
available
• Lazy Hogs !
Hund’s Rules cont.
1) One electron MUST occupy each orbital
BEFORE electrons are paired in the same
orbital.
2) Electrons added to subshell with the same
spin (+1/2, -1/2) so each orbital has one
electron.
Pauli Exclusion Principle
• Only 2 electrons occupy each orbital
• Electron spins MUST be opposite/paired when
2 electrons occupy the same orbital
– +1/2, -1/2

Using the periodic table-• Period numbers = principal quantum number
of valence shell electrons
• Subshells fill with electrons at different
regions within periodic table (s section, p
section)
Ex. 1 Nitrogen
Ex. 2 Mg
Practice
1) Boron
2) Copper
3) Sodium
The Name’s BOND…….
Chemical Bond
Diatomic Molecules
• Always exist in a chemical bond with another
atom, even if the atom is of the same element
• H2 , N2 , O2 , F2 , Cl2 , Br2 , I2
Chemical Bond
• Involves 2 atoms
• Mutual, electrical attraction between nuclei
and valence electrons of 2 atoms
IONIC BOND
bond formed between
two ions by the
transfer of electrons.
ELECTRON STEALING !
Ionic Bonds
• Between atoms of metals and nonmetals with
very different electronegativity
– Huge difference in electronegativity
– Electronegativity value > 1.7
• Bond formed by transfer of electrons
• Produce charged ions all
• Form between elements located on opposite
sides of the periodic table
COVALENT BOND
bond formed by the
sharing of electrons
Electron sharing!
Covalent Bond
• Between nonmetallic elements of similar
electronegativity.
– Small electronegativity difference <1.7
• Formed by sharing electron pairs,
• Formed between elements on the same side of
the periodic table.
NONPOLAR
COVALENT BONDS
when electrons are shared
equally, small electronegativity
difference
H2 or Cl2
POLAR COVALENT BONDS
when electrons are shared but
shared unequally
H 2O
Polar Covalent Bonds--More
• One atom “keeps” electrons closer to it
• Electrons tend to reside around one atom
more than the other atom
• Electrons still remain distributed between the
2 atoms--unequal
Identify Bond type, Draw Lewis
Structure to show bonding
NaCl, CO
METALLIC BOND
bond found in
metals; holds metal
atoms together
very strongly
Electrons
• Localized:
– Electrons hang out in a local area
– Restricted to an atom/ion or shared between
atoms
• Delocalized:
– Electrons not attached to a particular area
– Electrons are able to travel and move among
atoms
Metallic Bonds
• Formed between atoms of metallic elements
– Between atoms on left side of periodic table
– Electronegativities are approximately equal
• Electron cloud around atoms, delocalized
electrons
– Valence electrons shared among multiple atoms,
“neighbors”
• Electrons move through the whole metal, can
jump between energy levels to create
conduction bands.
Metallic Bond, A Sea of Electrons