Download ATOMIC STRUCTURE Text Book Chapters 2, 4, 5 OBJECTIVES

ATOMIC STRUCTURE
Text Book Chapters 2, 4, 5
OBJECTIVES
Matter is made up of particles whose properties determine the observable
characteristics of matter and its reactivity.
 Use models to describe the structure of an atom
 Relate experimental evidence to models of the atom
 Determine the number of protons or electrons in an atom or ion when given one
of these values
 Calculate the mass of an atom, the number of neutrons or the number of
protons, given the other two values
 Distinguish between ground state and excited state electron configurations
 Identify an element by comparing its bright-line spectrum to given spectra
 Distinguish between valence and non-valence electrons given an electron
configuration
 Draw a Lewis electron-dot structure of an atom
 Interpret and write isotopic notation
 Given an atomic mass, determine the most abundant isotope
 Calculate the atomic mass of an element, given the masses and ratios of naturally
occurring isotopes
Matter matters!
Chemistry = the study of the _______________________________________
Matter = anything that has ________________________________________
Elements = substance that ___________________________________________
by a chemical change; the simplest form of a pure substance
Example:
Atoms = the _________________ particle of an element that retains its identity
Molecule = a ___________ group of 2 or more atoms ______________ combined
Example:
Compounds = substances that contain ___________________________________
in a fixed proportion
Example:
Parts of the Atom
Subatomic particles:
Location
Protons
Neutrons
Electrons
Charge
Mass
Symbol
Atomic number = number of ________________ in the nucleus. Specific to an
element
Mass number = sum of _________________________________
((Note::::difference between atomic number and mass number is neutrons))
N
# neutrons = mass number - atomic number
ISOTOPES
Differences
Different number of ___________
(this will mean different _____!!)
Similarities
Same number of _____________
Same number of _____________
Examples of isotopes__________________________________________
Carbon–12 or 12C
____ Protons
____ Neutrons
____ Electrons
Carbon–14 or 14C
____ Protons
____ Neutrons
____ Electrons
BOTH ARE CARBON!!!!
Reminder: Isotopes are atoms of the ______________________ that have
different __________________________, and therefore have different
_____________________.
Complete the table:
Hydrogen Isotope
# of Protons
Protium
Deuterium
Tritium
# of Neutrons
Mass Number Symbol
AVERAGE ATOMIC MASS
Atomic Mass = the _____________ of all the isotopes in a sample of that element
Example: Look on the periodic table….what is the atomic mass of chlorine? _____
Chlorine has 2__________ that contribute to this atomic mass: _______________
Why is the average atomic mass not 36amu?
_______________________________________________________________
General Formula:
Average atomic mass = (percent/100) x (mass) + (percent/100) x (mass)……
Example: What is the average atomic mass of carbon if we are considering that
98.89% is carbon-12 and 1.108% is carbon-13?
Isotope name
Si-28
Si-29
Si-30
Isotope mass (amu)
percentage
92.23%
4.67%
3.10%
Find the average atomic mass of an atom of silicon.
CALCULATING ELECTRONS
1. Neutral : An atom has _______________ charge.
# protons = __________________
# electrons = _________________
Example: ___________
2. Ion : An atom with either a ________________________ (cation) or a
________________________ (anion)
# protons = __________________
# electrons = _________________________________
Example of a cation: ____________
Example of an anion: ____________
Calculate the number of electrons for the following:
1. Na
2. Na+1
3. Ca
4. Ca+2
5. O-2
6. O
7. Cl-1
HOW DID WE FIGURE ALL OF THIS OUT???
 2000 years ago Greeks thought matter was made up of particles of 4
elements:
________________, _______________, _______________, _______________,
Democritus 460-370BC said atoms were indivisible and unchangeable (but no one
listened to him)
 1600’s
Robert Boyle
He identifies gold and silver as elements – so they were not made of fire, earth,
air or water.
 1700’s
John Dalton
Dalton said that the basic unit of matter is a particle called the ____________
This is the basis for modern atomic theory
Dalton’s Atomic Theory
1.
2.
3.
4.
 1897
J.J. Thomson
Used a cathode ray to show atoms are made of smaller
parts – including the ___________________
Plum Pudding Model –
Atoms are neutral – so the electron (negative) and proton (positive) charges
must be equal in magnitude
 1909
Ernest Rutherford
Gold Foil Experiment – shot ______________ at a piece of
gold foil. Expected most to be deflected slightly, but saw
that most passed through and some were deflected
greatly (even bounced back!). So there must be a dense
positive core (nucleus!!)
Nucleus contains all positive charge and all the mass
 1913
Neils Bohr
Planetary Model –
Electrons only existed at specific circular orbits, and they
gain / lose energy when they shift from one energy level
Energy had previously just been viewed as a wave, while matter was made up
of particles. In the early 1900’s the WAVE MECHANICAL MODEL suggested a
dual nature model of the atom.
Same as the earlier models….
 Most of the atom is ____________________, with small positively charged
dense nucleus separated from electrons located outside
The difference is that
 Electrons do not stay in definite fixed orbits – rather they are probably
found in a region around the nucleus called an _________________
Picture of current model:
ELECTRON CONFIGURATION
 How the __________ are distributed throughout an atom’s energy levels.
While not contributing to its mass, this will determine an atom’s chemical
properties.
Valence Electrons = electrons in the ___________________ energy level
Kernel = everything except the __________________ electrons
Group =Vertical rows on the periodic table.
Name 3 elements in the same group_____________________
___________________________ = very similar chemical properties
The group number can be used to determine the number of valence electrons
Example:
Group 1 = ___ valence electron
Group ___ = 2 valence electrons
Group 14 = 4 valence electrons
Group 15 = ___ valence electrons
Lewis Dot Diagrams
Valence electrons are represented as dots around the element symbol.
Na
Cl
Li
Li1+
Ground State vs. Excited State
Ground State = When electrons occupy the lowest possible energy levels.
Maximum number of electrons allowed in each energy level is
2-8-18-32
When atoms are excited, energy is __________
Electrons jump to higher principal energy levels, which are higher _____________
Excited State = When __________ occupy ___________ energy levels before
occupying ______________________________________________________
Ground State:
Excited State:
2-8-1
2-7-2
Ground State:
Excited State:
2-8
Ground State:
Excited State:
2-8-1
Ground State:
Excited State:
2-4-1
BRIGHT LINE SPECTRUM
___________ absorbs a specific amount of energy.
Electron _________________________ energy level (__________________)
Electron releases ______________ that was ____________
Electron returns back to the ______ energy level
(__________________________)
Bright
Line
Spectrum
formed
This energy released is at a specific wavelength and so can predictably be seen in
the bright line spectrum for each element.
NAME: _______________________________________
ATOMIC STRUCTURE
BIG IDEAS
FORMULAS
REFERENCE TABLES
Atom
Atomic mass
Atomic mass unit
Atomic number
Average atomic mass formula
Bohr
Bright line spectrum
Cathode ray tube
Dalton
Dot diagrams
Electron
Electron configuration
Element
Excited state
Gold foil experiment
Ground state
Ion
Isotope
Mass number
Neutral
Neutron
Nucleus
Orbital
Planetary model
Plum pudding model
Proton
Rutherford
Thompson
Valence electrons
Wave-mechanical model (cloud model)