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

Group 1 – alkali metals – reactiveGroup 2 - alkaline earth metals – reactiveGroup 3-12 – transition metalsGroup 15 – nitrogen familyGroup 16 – oxygen family – reactiveGroup 17 – halogens – very reactiveGroup 18 – noble gases

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

- electrons are shared by two atoms

Naming Molecular1. Prefixes: (MEMORIZE)

Mono-1 tetra-4 hepta-7 deca-10di-2 penta-5 octa-8tri-3 hexa-6 non-9

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