Periodic Table and Bonding Notes Electromagnetic radiation – energy that travels through space as waves. Waves have three primary characteristics: 1. _______________ (- lambda) – distance between two consecutive peaks or troughs in a wave. Unit = meter 2. _______________ ( = nu) – indicates how many waves pass a given point per second. Unit = Hertz (Hz) 3. ________________ – velocity (c = speed of light = 3 x 10 8 m/sec) - indicates how fast a given peak moves in a unit of time c = Electromagnetic radiation (light) is divided into various classes according to wavelength. _______________________ – Light as waves – Light as photons (de Broglie) Photon/quantum – packet of energy – a “___________” of electromagnetic radiation Energy - (E – change in energy) – Unit Joules (J) Planck’s Constant – (h = 6.626 x 10 -34 J * s) E photon = hE photon = hc
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Periodic Table and Bonding Notes
Electromagnetic radiation – energy that travels through space as waves.Waves have three primary characteristics:1. _______________ (- lambda) – distance between two consecutive
peaks or troughs in a wave. Unit = meter2. _______________ ( = nu) – indicates how many waves pass a
given point per second. Unit = Hertz (Hz)3. ________________ – velocity (c = speed of light = 3 x 108 m/sec) -
indicates how fast a given peak moves in a unit of time
c = Electromagnetic radiation (light) is divided into various classes according to wavelength.
_______________________ – Light as waves – Light as photons (de Broglie)
Photon/quantum – packet of energy – a “___________” of electromagnetic radiation
Energy - (E – change in energy) – Unit Joules (J)Planck’s Constant – (h = 6.626 x 10-34 J * s)
Ephoton = h Ephoton = hc
Ex: What is the wavelength of light with a frequency of 6.5 x 1014 Hz? What is the change in Energy of the photon?
So with light waves, you can convert between wavelength, frequency, and energy with two equations: = c E = h And two constants:
c = 3 * 108 m/s h = 6.626 * 10-34 J sIn the visible part of the spectrum, different colors correspond to
different frequencies, wavelengths and energies. Blue light has a ____________ wavelength, __________ frequency and _____________ energy. Red light has a ___________ wavelength, __________ frequency, and ___________ energy.
__________________ – atom with excess energy
__________________ – lowest possible energy state
Wavelengths of light carry different amounts of energy per photon
Only certain types of photons are produced (see only certain colors)
_______________ – only certain energy levels (and therefore colors) are allowed
Emission Spectrum – bright lines on a dark background. Produced as excited electrons return to a ground state – as in flame tests.
Absorption Spectrum – dark lines in a continuous spectrum. Produced as electrons absorb energy to move into an excited state, only certain allowable transitions can be made. Energy absorbed corresponds to the increase in potential energy needed to move the electron into allowed higher energy levels. The frequencies absorbed by each substance are unique.
Intensity
Color
Bohr Model – suggested that electrons move around the nucleus in circular orbits
Only Correct for Hydrogen
Wave Mechanical Model – Described by orbitals gives no information about when the electron occupies a certain
point in space or how it moves *aka – Heisenberg’s uncertainty principle
Parts of the Wave Mechanical Model1. Principle Energy Level (n) – energy level designated by numbers 1-7.
called principle quantum numbers
2. Sublevel – exist within each principle energy level the energy within an energy level is slightly different each electron in a given sublevel has the same energy lowest sublevel = s, then p, then d, then f
3. Orbital – region within a sublevel or energy level where electrons can be found
s sublevel – 1 orbitalp sublevel – 3 orbitalsd sublevel – 5 orbitalsf sublevel – 7 orbitals- ** No more than two electrons can occupy an orbital**- an orbital can be empty, half-filled, filled
Electron Configuration – arrangement of the electrons among the various orbitals of the atom
1
2
3
4
5
6
7
lectronegativitylectronegativity
Ex:
Shapes of orbitals- All s orbitals are spherical as the principle energy level increases
the diameter increases.
- All p orbitals are dumbbell shaped – all have the same size and shape within an energy level
- All d orbitals are flower (clover) shaped and a donut – all have the same size and shape within an energy level
Electron Spin Spin – motion that resembles earth rotating on its axis– clockwise or counterclockwise
Pauli Exclusion Principle – two electrons in the same orbital must have opposite spins
Hund’s Rule – All orbitals within a sublevel must contain at least one electron before any orbital can have two
Orbital Diagram – describes the placement of electrons in orbitals- use arrows to represent electrons with spin- line represents orbital (s=1, p=3, d=5, f=7)
____ full ____ half-full ____ empty Ex:
Noble Gas Configuration – Shorthand configuration that substitutes a noble gas for electronsEx:
Valence Electrons – Electrons in the outermost (highest) principle energy level in an atom
Core Electrons – innermost electrons – not involved in bonding
Valence Configuration – shows just the valence electronsEx
Periodic Table
Dimitri Mendeleev-1869- developed the first version of the periodic table.He expressed the regularities as a periodic function of the
_______________.
Henry Moseley- revised Mendeleev periodic table by describing regularities in physical and chemical properties as periodic functions of the ____________.
Groups (family) – vertical columnElements with similar ________________________ configurations
Periods – horizontal rowsPeriod number corresponds to the ____________________________ of
valence electrons
Periodic Trends1. Atomic Size (radius)
Increases – down a groupDecreases – across a period
Size of ionsCation Ca+2/Ca Ca larger because Ca+2 lost 2 electronsAnion S-2/S S-2 larger because S-2 gained 2 electrons
2. Ionization Energy – energy required to remove an electron from an individual atom in a gas phase M(g) M+(g) + e-
• Metals lose electrons to non-metals so relatively low energy is needed
• High ionization energy means an electron is hard to removeDecreases – down a groupIncreases - across a period
3. Electron Affinity – Electron affinity is the energy involved when an electron is added to a gaseous atom.
• Negative values of energy mean that energy was released during the process. Atoms with negative values of electron affinity have a very strong attraction for electrons.
• Positive values of electron affinity have very little attraction for electrons.(energy involved in negative ions)
Decreases – down a groupIncreases - across a period
4. Electronegativity is the tendency of an atom to draw electrons to itself when in a covalent bond. Consequently, the trends are the same as for electron affinity. The atoms with the highest electronegativity are fluorine, then oxygen, then nitrogen. It is also important to know that the electronegativity of hydrogen is slightly less than that of carbon.
Decreases – down a groupIncreases - across a period
5. Metallic Character Increases – down a groupDecreases – across a period
Summary of Trends
Chemical Bonding Notes
Bond- force that holds groups of two or more atoms together and makes them
function as a unit
bond energy- energy required to __________ the bond (tells the bond strength)
Ionic bonding- between ionic compounds which contain a ____________ and a ____________________
Atoms that lose electrons relatively easily react with an atom that has a high affinity for electrons
Transfer of electrons
Covalent bonding- between two nonmetals Electrons are ______________ by nuclei
Polar Covalent bonding- ________________ sharing of electrons positive end attracted to the negative end (delta) indicates partial charge
_______________________ - relative ability of an atom in a molecule to attract shared electrons to itself
The higher the atom’s electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bonds
Increases – across a period Decreases- down a group
Electronegativity difference Bond type
Covalent character
Ionic character
Zero Covalent Decreases IncreasesIntermediate Polar covalent Decreases Increases
Large Ionic Decreases increases
Ex. List the following in order of increasing polarity.H-H, O-H, Cl-H, S-H, F-H
Dipole moment- has a center of positive charge and a center of negative charge
Represented by an arrow Arrow points toward the negative charge
Chemical Formula – type of notation made with ________________ and chemical symbols
- indicates the composition of a compound- indicates the number of atoms in one molecule
Molecule – covalently bonded collection of two or more atoms of the same element or different elements
- monatomic molecule – one atom molecules- diatomic molecule – two atom molecules (seven) MEMORIZE
Br, I, N, Cl, H, O, F
MetalsLocation: ____________ side of Periodic TableProperties: Ductile – drawn into wires
Malleable – hammered into sheetsMetallic Luster – ________________Good Conductors of Heat and Electricity
NonmetalsLocation: ___________ side of Periodic Table
Properties: BrittleLack Luster – not shiny
Poor Conductors of Heat and ElectricitySemi-metalsLocation: Along Stair-step
Properties: Have properties of metals and nonmetals- also called METALLOIDS- Si, Ge, As, Sb, Te, Po, At
Molecular NomenclatureMolecular Compounds (molecules) – compounds made from two nonmetals
2. prefixes are used with both the first named and second named element. Exception: mono- is not used on the first word
3. second word ends in –ide4. If a two syllable prefix ends in a vowel, the vowel is dropped
before the prefix is attached to a word beginning with a vowel
Writing molecular formulasTranslate prefixes
Examples:
Valence electrons are used in bonding.
N2O
Si8O5
NH3
P3I10
dihydrogen monoxide
tetrasulfur hexachloride
carbon monoxide
carbon dioxide
Stable elements want to achieve 8 electrons similar to the noble gases
If it’s a metal it wants to achieve the configuration for the noble gas before.
If it’s a nonmetal it wants to achieve the configuration for the noble gas after.
Lewis Structure- representation of a molecule Shows how the valence electrons are arranged among the
atoms in the molecule.For an element:
For a compound:
For a molecule:
Duet rule- only two electrons in the full shell
Octet rule- surrounded by eight electrons
Bonding pair- electrons shared with other atom
Lone pair or unshared pair- not involved in bonding
5 Steps for Covalently Bonded Lewis Structures1. Find the total number of valence electrons.2. Calculate the number of “needed” electrons to give each atom 8
electrons, except for H which wants 2.3. Subtract valence electrons from the “needed” electrons. This is the
number of bonding electrons.4. Divide the bonding electrons by 2, to find the number of bonds.5. Subtract the bonding electrons from the valence electrons to find
the non-bonding electrons or lone pairs.6. Choose a central atom and assemble the pieces to make all atoms
involved stable.
Ex. GeBr4
Single bond- involves two atoms sharing one pairDouble bond- involves tow atoms sharing two pairsTriple bond- involves two atoms sharing three pairs
Ex. CH4 C2H4 C2H2
Resonance- more than one Lewis structure can be drawn for the molecule
Ex. CO2
Exceptions to the Octet Rule1. boron and beryllium - tend to be electron deficient
boron can hold 6 total electrons beryllium can hold 4 total electrons
ex. BF3 BeH2
2. Electrons are small spinning electric charges that create magnetic fields
Diamagnetic - substances which have paired electrons that cancel out the magnetic field
Paramagnetic - substances the have one or more unpaired electrons that show great attraction to the magnetic field
Ex. O2 PH3
3. Odd number of electrons You cannot write electron dot structures that fulfill the octet rule,
when the total number of valence electrons is oddEx. NO
4. Expanded Octet- expand the valence shell to include more than 8 electrons Phosphorus and sulfur can expand to include 10 or 12 electrons You will know you have an expanded octet when you don’t have
enough bonds for the atoms presentEx. SF6
StructureMolecular (geometric) structure- three-dimensional arrangement of the atoms in a molecule
VSEPR model- valence shell electron pair repulsion Lone pairs of electrons like to be as far away from each other
as possible Double and triple bonds “act” like a single shared pair for
shape.
1. Linear- two pairs of electrons are present around an atom One total pair – one shared pair Two total pairs – two shared pairs Bond angle = 180
Ex. BeCl2
2. Bent Four total pairs Two shared pairs and two unshared
pairs Bond angle = 104.5
Ex. H2O
3. Trigonal planar - whenever three pairs of electrons are present they should be placed at the corners of a triangle
Three total pairs Three shared pairs Bond angle = 120
Ex. BCl3
4. Tetrahedral Four total pairs Four shared pairs no unshared pairs Bond angle = 109.5
Ex. CCl4
5. Trigonal pyramid Four total pairs
Three shared pairs and one unshared pair Bond angle = 107Ex. NH3